Chemistry Unit 2

During chemical reactions there is often a difference between the amount of energy taken in to break the bonds and the amount of energy given out in the formation of new bonds.

Exothermic reactions give out energy overall. Substances involved get hotter and the chemicals lose energy, this energy is gained by the surroundings. This means the enthalpy change is always negative

Endothermic reactions take in more energy than they release. Substances involved get colder and the energy that is lost by the surroundings is gained by the chemicals. This means the enthalpy change is always positive

Standard enthalpy change
These are used so that different enthalpy changes can be compared. It means the conditions must be constant.

• 1 mole of a substance is reacted
• All solutions have a concentration of 1dm^-3
• All gases have a pressure of 100kPa
• Temperature must be stated, usually 298K
• Elements must be in their standard states 

Many reactions are reversible. If a reversible reaction takes p[lace in a closed system then an equilibrium can be reached. At equilibrium there is a balance between the reactants and the products. At equilibrium the rate of the forward reaction is equal to the rate of backward reaction. 

Le Chatelier's principle states that when the conditions of a dynamic equilibrium are changed then the position of equilibrium will shift to minimize the change.

As the temperature is increased the position of equilibrium shifts in the endothermic direction.

As the pressure is increased the position of equilibria shifts towards the side with the least gas molecules.

As the concentration of one of the reactants is increased the position of equilibrium shifts towards the right so more product will be made.

The Haber Process
Ammonia is produced via a reversible reaction. The forward reaction is exothermic

N2(g) + 3H2(g) <==> 2NH3(g)

An iron catalyst is used and has no effect on the yield, catalysts provide an alternative reaction pathway with a lower activation energy. As the forward reaction is exothermic, increasing the temperature increases the rate but decreases the yield. Therefore a compromise temperature is used. Increasing the pressure increases the yield. 

The contact process
Sulfuric acid is produced via a reversible reaction. The forward reaction is exothermic.

SO2(g) + 0.5O2(g) <==> SO3(g) 

A vanadium catalyst is used and has no effect on yield. Increasing the temperature increases the rate and decreases the yield, again a compromise temperature is used. Increasing pressure increase the yield, though in practice high pressure isn't used as it is too expensive.

The hydration of ethene
Ethanol can be made by the hydration of ethene.
Ethene + steam <==> ethanol vapour
The forwards reaction is exothermic.
A phosphoric acid catalyst is used. Increasing the temperature increases the rate of reaction but decreases the yield of ethanol.

Production of methanol
Methanol can be made by reacting carbon monoxide with hydrogen
Carbon monoxide + hydrogen <==> methanol
If an equilibrium mixture of carbon monoxide, hydrogen and methanol is placed in a fuel tank, as the hydrogen is used more hydrogen will be produced.

Carbon neutral
The production of methanol is carbon-neutral if the hydrogen is produced from a renewable source. Also, the carbon monoxide can be made from plant waste. 

Fuels
Both are liquid fuels at room temperature. These are easier to transfer than gas and take up a smaller volume. They both burn cleanly releasing very little carbon monoxide.
Ethanol made by fermentation is a renewable fuel 

• Oxidation is the loss of electrons
• Reduction is the gain of electrons

If one species loses electrons another must gain them, meaning that oxidation and reduction must always occur together, redox reactions

Agents
Oxidizing agents are species that oxidize other substances by removing electrons from them. They are themselves reduced. 
Common examples include: oxygen, chlorine, potassium dichromate and potassium manganate.
Reducing agents are species that reduce other substances by adding electrons to them. They are themselves oxidized. 
Common examples include: group 1 & 2 metals, hydrogen, carbon and carbon monoxide.

Any pure element has an oxidation state of zero

Any monatomic ion has an oxidation state equal to the charge on the ion

In compounds, group 1 metals have an oxidation state of 1+

In compounds, group 2 metals have an oxidation state of 2+

In compounds, fluorine always has an oxidation state of 1-

In compounds, hydrogen has an oxidation state of 1+ unless it is in a metal hydride, then it is 1-

In compounds, oxygen has an oxidation state of 2- unless it is with fluorine when it has an oxidation state of 2+, or it is part of peroxide when it has an oxidation state of 1-

If a molecule is neutral overall, then the sum of the oxidation states must be zero

When a molecule has an overall charge, the sum of the oxidation states must be equal to the charge

Halogens exist as diatomic molecules, this means that there a re strong covalent bonds within the halogen molecules but only very weak vdW forces of attraction between halogen molecules.
The strength of the vdWs depends on the number of electrons, this increases down the group.

Trends

Electronegativity: down the group the electronegativity decreases 

Boiling point: down the group the boiling point increases

Oxidizing ability: down the group the oxidizing ability decreases

Reducing ability: (halide ions) down the group the reducing ability increases 

We can identify halide ions using acidified silver nitrate solution
Fluoride: no precipitate
Chloride: white precipitate
Bromide: cream precipitate
Iodide: yellow precipitate

Then add dilute ammonia
AgCl: redissolves
AgBr: no visible change
AgI: no visible change

Then add concentrated ammonia 
AgBr: redissolves
AgI: no visible change 

Mg
Ca
Sr
Ba 

Trends

Atomic radius: increases down the group

First ionization energy: decreases down the group

Reducing ability: increases down the group

Melting point: increases down the group

Reactivity with water: increases down the group 

The solubility of group 2 hydroxides increases down the group

Mg(OH)2: (in solution) milk of magnesia, it's used to treat indigestion and heartburn. It neutralizes excess hydrochloric acid
Ca(OH)2: neutralizes soil ('lime'). Formed from calcium carbonate. 

The solubility of group 2 sulfates decreases down the group

MgSO4.7H2O: Epsom salts; used in fireproofing, artificial snow, laxative and bath salts 
2CaSO4.H2O: Plaster of Paris; used in plaster casts and plasterboard 

BaSO4: opaque to X-rays so is regularly used in medicine, though it is toxic 

Extraction of iron: (in blast furnace)

Carbon monoxide reduces the iron oxide (found in haematite) to iron and carbon dioxide. 

Extraction of manganese: 
Manganese oxide (found in pyrolusite) is reduced, first by carbon monoxide, then by carbon 

Extraction of copper:
Copper oxide (found in chalcopyrite) is reduced by carbon

Extraction of aluminium: (electrolysis)