5b) Crude Oil

Crude Oil is a dark, smelly liquid which is a mixture of lots of different chemical compounds. A mixture contains two or more elements/ compounds that aren't chemically bond together.

Fractional Distillation:
-separates liquids with different boiling points

• Oil heated as it enters fractioning column 
• Chemicals in oil evaporate 
• Those with lowest boiling points evaporate first 
• They then condense to liquids 
• Lowest boiling point chemicals move further up the fractioning column 
• Liquid fractions drained off the column for use 

Fraction	Boiling Range	Viscosity	Use	Chain length
Refinery gases	lowest	lowest	bottled gases	Shortest C1 - C4
Gasoline			Petrol for cars	C5 - C9
Kerosine			Jet fuel	C12 - C15
Diesel			Fuel for cars/lorries	C15 - C20
Fuel oil			Fuel for ships	C21 upwards
Bitumen	highest	highest	Road buildings/ roofs	Longest around C30

Incomplete combustion can produce carbon monoxide which is dangerous because it reduces the capacity of the blood to produce oxygen.

Sulfur dioxide and nitrogen come from burning fuel. The sulfur dioxide comes from sulfur impurities in the fossil fuels. Nitrogen oxides are created when the temperature is high enough for the nitrogen and oxygen in the air to react. This often happens in car engines. Nitrogen monoxide and nitrogen dioxide are included in nitrogen oxide. 

Pollutants:

Acid rain is caused by sulfur dioxide and nitrogen oxides 

When sulfur dioxide mixes with the clouds it forms dilute sulfuric acid, which is much more acidic

2SO2(g) + O2(g) + 2H2O ---> 2H2SO4(aq) 

Nitrogen oxides can also form nitric acid in clouds, this rain is called acid rain. 

Acid rain causes lakes to become acidic and many plants and animals die as a result and it also kills trees and damages limestone buildings and ruins stone statues.

Cracking

Converts long chain alkanes in the heavier freactions into shorter chains that are currently in more demand. Also as a result an alkene is also produced which is used to make polymers.

Conditions for cracking

-In industry vapourised hydrocarbons are passed over a powered catalyst at about 600C - 700C. Silica (SiO2) and alumina (Al2O3) are used as catalysts.

1. heat the paraffin (or another alkane), after a few seconds move the bunsen burner to heat the catalyst. Alternate between the two until paraffin vapourises and catalyst glows red
2. heated paraffin cracks as it passes over catalyst 
3. small alkanes collected at end of boiling tube and alkene gases travel down the delivery tube 
4. Alkenes are collected through water using a gas jar

5a) Extraction and Uses of metals

Methods of extraction are linked to the order of reactivity

Only metals that are less reactive than carbon can be extracted by a reduction reaction with carbon- this is done by heating the ore with carbon monoxide. This is because more reactive elements form compounds more readily.

Metals more reactive than carbon have to be extracted using electrolysis. It uses electricity to separate the metal from the other elements in the compound.

Aluminium extraction


Electroylsis is used to remove aluminium from its ore. A catalyst- Cryolite is used to lower the temperature and cost as Al2O3 has a very high boiling point of 2000C and so melting it would be very expensive. It is dissolved into molten cryolite which brings the temperature down to 900C, the electrodes are made of graphite which is a good conductor of electricity.



At the negative electrode (Cathode): Al³⁺ + 3e^- --> Al
At the positive electrode (Anode): 2O₂- --> O₂ + 4e^-

Extracting Iron:

(Blast Furnace) 

Raw materials:

1. Iron ore -----> iron
2. Coke -----> almost pure carbon; reducing iron oxide to iron metal
3. Limestone -----> take away impurities from slag 
4. Air ------> allows the coke to burn 

Reducing Iron ore to Iron:

1. Hot air blasted into the furnace -----> coke burns faster than normal. Raises to 1500C.  
2. Coke burns and produces carbon dioxide:
• Carbon and Oxygen ------> Carbon Dioxide 
• C            +        O2    ------> CO2 
3. Carbon monoxide 
• Carbon dioxide and Carbon -----> Carbon monoxide 
•          CO2          +         C     ----->          2CO 
4. Carbon monoxide then reduces the iron ore to iron
• Carbon monoxide + Iron(II) oxide ------> Carbon dioxide + Iron
•          3CO             +       Fe2O3       ------>        3CO2        +  2Fe 
5. Iron is molten at temperature also dense so goes to the bottom where it's tapped off 

Removing impurities: 

Slag forms as CaO from limestone reacts with rocks (SiO2) 

CaCO3 -----> CaO + CO2 

CaO + SiO2 ------> CaSiO3 (molten slag) 

• Coolen to solid and used for:
• road building 
• fertiliser 

Properties of both Iron and Aluminium

• both dense and shiny 
• have high melting points 
• strong and hard to break 
• are malleable 
• good conductors of heat and electricity 

Uses of Iron

• steel and stainless steel 
• buildings
• car radio 
• railings

Uses of Aluminium

• doesn't corrode so is useful for products that come in contact with water such as drinks cans
• Used in bicycle frames and aeroplanes when the weight has to be very specific 

Section 5

Chemistry in Industry

4d) Equilibria

Some reaction are reversible which means that they have the same rate backwards and forwards
That means that there is no overall change: reactants --> products    products --> reactants

This can only take place in a CLOSED system or it would escape. Initially there are no products when you first start to mix the reactants. After a while there will be products; but still more reactants. With time the reactants go down and the products go up and we reach the SAME RATE, this is called DYNAMIC EQUILIBRIUM.
The position of equilibrium is not always at the half way point therefore there could be a position where there is more reactants than products.

Dehydration of hydrated copper (II) sulfate:
hydrated copper(II) sulfate + heat   anhydrous copper(II) sulfate + water

Heat makes the reaction go forwards (endothermic), whilst cold water makes it go backwards (exothermic) 

Effect of heat on ammonium chloride:

ammonium chlorideammonia + hydrogen chloride

Ammonium chloride decomposes when it is heated so the forward reaction is endothermic, whilst the backwards reaction is exothermic 

Factors which influence equilibrium:

• Temperature increase
• moves the equilibrium to the right hand side, in the direction that produces the fewer molecules/moles of gas on the RHS  
• Pressure increase
• moves equilibrium to the right hand side, moves into the direction that absorbs heat energy e.g. endothermic reactions 
• Concentration 
• Catalyst to get products or reactants faster 

 

LHS              RHS 

A + B  C + D 

4c) Rates of Reaction

RATE - ^{AMOUNT REACTION}⁄<sub>TIME

Depends on four things:</sub>

1. Temperature 
2. Concentration
3. Catalyst 
4. Size of particles (Surface area) 

The Collision Theory:
Higher Temperature: particles move faster, collide more often and with more force. If the temp. is high there will be enough energy to overcome Ea.

Higher Concentration: more particles in more conc., so they collide more frequently

Large surface area: smaller pieces increase sa, more area to work on, collisions more frequent

Catalyst: increase the number of successful collisions by lowering the Ea.

Activation energy or Ea is the amount of energy needed to start a reaction
(higher temp. increases reaction)

4b) Energetics

Exothermic reactions:
Energy is transferred from the reacting mixture to the surroundings and temperature of surroundings increases

Endothermic:
Energy is transferred from the surroundings to the reaction mixture and temperature of surroundings decreases

Copper cup experiment is a simple calorimetry experiment:

Paper Two - Molar Enthalpy:
Measure temperature change of the experiment
Then use this calculation - Energy, Q = mc∆T                  m= mass of surroundings (water or solution) ∆T = change in temp.
                                           or

                                           ∆H = Q/n                                  n = number of moles    sign - exo and + endo 

∆H is the change in heat energy 

Q is measured in Joules 

Simple Energy Level Diagram:

Exothermic

Endothermic 

Ea or activation energy is needed to break or make bonds

Bonds:

Breaking bonds puts energy in - endothermic, 

whilst making new bonds releases energy - exothermic

Paper Two - Using average bond energies to calculate the enthalpy change during a simple chemical equation:

You are given the reaction and table of data 

Then with the information work out the broken bonds (A) and the bonds you make (B)  

Then subtract them to get the ∆H 

4a) Acids and Alkalis

Indicators:

1. Universal indicator is a very useful combination of dyes which give one of the colours on the pH scale
2. Litmus paper tests whether a solution is acidic or alkaline because it changes colour at pH 7.  
1. Red in acidic solutions 
2. Blue in alkaline solutions 
3. Purple in neutral solutions 
3. Phenolphthalein will change from colourless in acidic solutions to bright pink in alkaline solutions 
4. Methyl Orange changes from red in acidic solutions to yellow in alkaline solutions. 

pH Scale:

+----+---------------------------+---------+--------------------------------+
|    |           Acids           | Neutral |            Alkaline            |
+----+---------------------------+---------+--------------------------------+
| pH | 0 | 1 | 2 | 3 | 4 | 5 | 6 |    7    | 8 | 9 | 10 | 11 | 12 | 13 | 14 |
+----+---+---+---+---+---+---+---+---------+---+---+----+----+----+----+----+

Acids;

• below pH scale 
• produce hydrogen ions, H+ when dissolved in water 
• acids + metal ----> salt and hydrogen 
• acids + metal oxides ----> salt and water 
• acids + metal carbonates ----> salt and water and carbon dioxide 

Common acids and salts produced

-hydrochloric acid to chloride

-nitric acid to nitrates 

-ethanoic acid to ethanoates 

Alkalis:

• above pH 7 
• produce OH- (hydroxide) ion when dissolved into water
• alkalis + acids ----> salt and water 

Common alkalis 

-sodium hydroxide 

-ammonia solution 

General rules for predicting the solubility of salt in water:

• All common sodium, potassium and ammonium salts are soluble 
• All nitrates are soluble 
• Common chlorides are soluble, except silver chloride 
• Common sulfates are soluble, except those of barium and calcium
• Common carbonates are insoluble, except those of sodium, potassium and ammonium 

Preparing soluble salts from acids and insoluble bases:

• Add excess solid to ensure that all the acid has reacted 
• Filter it 
• Evaporate half water 
• Leave to crystallise 

Insoluble salts using precipitation reactions:

• Mix the two solutions which contain the ions you need 
• filter precipitate
• wash 
• dry

Acid-Alkais titration:

• Using a pipette and pipette filler, add some alkali to a conical flask, along with two or three drops of indicator 
• Fill a burette with acid 
• Using the burette, add the acid to the alkali a bit at a time- giving the conical flask a regular swirl. Go especially slowly when you reach the end point of the colour change of the indiactor
• The indicator changes colour when all the alkali has been neutralise 
• Record the volume of acid used to neutralise the alkali, it's best to repeat the process several times to make sure you have reliable results.