Condensation Polymerisation-Nylon

Note: This is for SINGLE SCIENCE.

How to make nylon: (he uses a diamine which is one of the monomers, but doesn't use a dicarboxylic acid so just beware of that. the video's just essentially to show you what it looks like to make nylon, stuff about the monomers are below)

Reviewing addition polymerisation vs. condensation polymerisation. This is a really good video to show you the difference and to explain it. It may not be about nylon, but the concept is essentially the same and it also loses a water molecule in the process, which is the same as when nylon is produced. 

5.17 recall that nylon is a condensation polymer

5.18 understand that the formation of a condensation polymer is accompanied by the release of a small molecule such as water or hydrogen chloride

Condensation polymers are basically polymers formed through a condensation reaction, where the monomers react and a polymer is produced, and a small molecule such as water is also produced as a by-product of the reaction. With condensation polymers the monomers can be different.

(As opposed to addition polymers which are produced by the reaction of unsaturated monomers. See polymerization post…)

5.19 recall the types of monomers used in the manufacture of nylon

http://ibchem.com/drop/nylon2.htm

Nylon is made by condensation polymerization from the monomers dicarboxylic acid and diamine.

NH₂ is the amine group. So a diamine has two of these, one at each end: 

 

COOH is the carboxyl group. So a dicarboxylic acid has two of these, one at each end:
  

I believe there are loads of types of nylon and different monomers to make them but this is what I've learned: 

The polymer made from these two six-carbon monomers is known as nylon 6,6. (Nylon products include parachutes and ropes.)
The diamine to form nylon here is 1,6-diaminohexane. Hexane=6, it has 6 carbons as you can see in the diagram below. The adipic acid is basically 1,6 hexane dicarboxylic acid. It also has 6 carbons if you count the carbons from the COOH groups. 

Instead of having to draw out all the carbons, you can replace them with 'R' and just show the functional groups that are reacting: 

So you see that the 'OH' on one end of the COOH group forms water with the 'H' from one of the amine groups. --> H₂O 

And the 'CO' left bonds with the 'NH'. 

This keeps happening at both ends so they form a long chain polymer. So you see in the above diagram it says 'carboxyl group for further reaction'? The 'OH' there will react with the 'H' from another amine group, remember they are diamines so there are amine groups on both sides. And the CO left will bond with the 'NH' left, and it'll just continue until there are no more reactants left. 

5.20 draw the structure of nylon in block diagram format.

You can replace the carbons in the middle with blocks like this, it would be good if the blocks were different shapes unlike below, just so it's clearer for you, but it's up to you: 

Reactivity Series

Section 2: f) Reactivity series

Seems boring at first but worth watching, and the teacher's awesome, you'll see. (this links to spec 2.30)

sodium + water in a 40 gallon trash can (y) 

love this one, 'it's coming for you..'

2.30 recall that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

Mnemonic	Element	Symbol	Reactivity
Please	Potassium	K	As you can see these metals (excluding carbon) are above hydrogen in the reactivity series so they react with acids and displace hydrogen gas. Metal + acid à metal salt + hydrogen
Send	Sodium	Na
Little	Lithium	Li
Charles	Calcium	Ca
McClean	Magnesium	Mg
A	Aluminium	Al
Common	Carbon	C
Zebra	Zinc	Zn
If	Iron	Fe
The	Tin	Sn
Lame	Lead	Pb
Horse	Hydrogen	H	H+ ions are responsible for acidic properties.
Can’t	Copper	Cu	These elements are below hydrogen so they do not react with acids. (Acids contain H+ ions) Exception: Copper reacts with concentrated nitric acid, the nitrate ions oxidize copper. But that’s not really important)
Munch	Mercury	Hg
Some	Silver	Ag
Grass	Gold	Au
Properly	Platinum	Pt

2.31 describe how reactions with water and dilute acids can be used to deduce the following

order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron, and

copper

Reactions with cold water: 

Basically, you will see that the higher up the metal in the reactivity series, the more vigorous the reaction. For example, the reaction of some alkali metals and water:

Alkali metal	Hydroxide solution produced	Gas produced	Rate of gas produced
Potassium	Potassium hydroxide	Hydrogen	Very vigorous
Sodium	Sodium hydroxide	Hydrogen	Vigorous
Lithium	Lithium hydroxide	Hydrogen	Fairly vigorous

They love asking about sodium. 

"Describe what happens when sodium is added to water" and stuff.. 

Well when sodium is added to water, it reacts very quickly and vigorously. It's an exothermic reaction and the heat produced causes the sodium to melt. The molten sodium darts around the water surface and a yellow flame is seen. You may see a bit of fizzing/bubbling (effervescence) as hydrogen is evolved. 

Remember MM-FF. Melts, moves, floats, fizzes

With Calcium, it reacts gently with cold water you may see some bubbles and calcium hydroxide is formed, or better known to some as limewater. So you will see a white precipitate forming as hydroxides are actually insoluble unless it's an alkali metal hydroxide. (All alkali metal salts and hydroxides are soluble.) 

Magnesium reacts slowly in water but reacts vigorously with steam. The reason why magnesium doesn't really react with cold water is that it becomes coated with magnesium hydroxide, which is insoluble, so it prevents water coming into contact with the magnesium. Magnesium also burns with a bright white flame and white magnesium oxide ash is formed. 

Zinc and iron don't react with cold water, but they react with steam to form oxides. Neither metal burns like magnesium. 

Copper doesn't react with water or steam as it is below hydrogen in the reactivity series. 

Reactions with acids: 

I'd say potassium, sodium, lithium and calcium are probably too reactive to react with dilute acids and would be quite dangerous. Too reactive to add safety to acids. They're already pretty violent with water. Basically the reaction would be exothermic and a lot of heat is produced, hence the hydrogen evolved could ignite and catch fire. 

Metal + acid à metal salt + hydrogen

You can tell the metal's position in the reactivity series by seeing how many bubbles are formed, or how fast. 

You can also: add a small piece of the metal to some cold water. If there is any rapid reaction, then the metal must be above magnesium in the reactivity series. If there isn't any reaction, add a small amount of the metal to some dilute hydrochloric acid or dilute sulphuric acid. If there isn't any reaction in the cold, warm/heat it carefully. 

If there's still no reaction, the metal is probably below hydrogen in the reactivity series. If there is a reaction, then it is somewhere between magnesium and hydrogen. 

2.32 deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

Any metal higher in the reactivity series will displace one lower down from its compound. So for example a reaction with magnesium and copper (II) oxide will result in the magnesium displacing (pushing out) the copper from its oxide, so the magnesium basically replaces it. 

Magnesium + copper (II) oxide à magnesium oxide + copper

Mg (s) + CuO (s) à MgO (s) + Cu (s)

It's the same thing with metals and a solution of their salt. The more reactive metal will displace a less reactive metal. For example, the reaction between zinc and copper (II) sulphate solution:

The copper is displaced by the more reactive zinc. The blue colour of the copper (II) sulphate solution fades as colourless zinc sulphate solution is formed. 

Zinc + copper sulphate à zinc sulphate + copper

Zn (s) + CuSO₄ (aq) à ZnSO₄ (aq) + Cu (s) 

2.33 understand oxidation and reduction as the addition and removal of oxygen respectively

Oxidation could mean the addition of oxygen, and reduction could mean the removal of it, but also remember that it can be about electrons: 

OILRIG = Oxidation is Loss, Reduction is Gain

So if something has lost electrons, it is oxidised.

Likewise if something has gained electrons, it has been reduced. 

2.34 understand the terms redox, oxidising agent and reducing agent

A redox reaction is a reaction in which both reduction and oxidation are occurring. They always go together. 

An oxidising agent is a substance that causes another substance to be oxidised. So it causes something else to lose electrons, and gains these electrons itself. So the oxidising agent itself is reduced. *This confuses people!! Remember that oxidising agent doesn't get oxidised, don't let the name fool you.

An example of good oxidising agents are the halogens. Especially fluorine, which is super reactive. They only need to gain one electron to get a full outer shell so they easily oxidise other elements, such as the alkali metals which only need to lose one electron too. A common example is sodium chloride-NaCl--your common table salt. 

To oxidise something can also involve oxygen, where oxygen is added to a substance. (See previous specification point)

A reducing agent is a substance that reduces something else. So it causes the substance to gain electrons, by losing electrons itself. So the reducing agent is said to be oxidised. It can also be taken as the reducing agent takes away oxygen from the other substance, such as: 

Magnesium + copper (II) oxide à magnesium oxide + copper

So here the magnesium is the reducing agent, whilst the copper (II) oxide is the oxidising agent. 

2.35 recall the conditions under which iron rusts

Iron rusts in the presence of oxygen and water. Rusting is accelerated in the presence of electrolytes such as salt. 

Note: Many metals corrode, but it is only the corrosion of iron that is referred to as rusting. 

2.36 describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

Obviously to prevent rusting you need to keep oxygen and water away from the iron. You can do this by painting it, or coating it in oil/grease, or covering it with plastic. But once the coating is broken, the iron will rust. 

Coating the iron with a metal below it in the reactivity series (such as tin) is just a barrier method. Once the layer of tin on the iron is scratched, a tin can, for example, will rust very quickly. This is because the iron is more reactive than the tin and the tin won't prevent it. 

2.37 understand the sacrificial protection of iron in terms of the reactivity series.

Galvanised iron is iron that is coated with a layer of zinc. It serves as a barrier to air and water. But unlike tin, if it is scratched, the iron still doesn't rust. This is because zinc is more reactive than iron, and so corrodes instead of the iron. So the zinc is 'sacrificed' for the iron.

Galvanising is the term used only when iron/steel is coated with a protective layer of zinc, with other metals, it is sacrificial protection. 

Electrolysis

Electrolysis

1.47 understand an electric current as a flow of electrons or ions

Current is basically a flow of charged particles and in a metal wire, it is a flow of electrons. These have a negative 1 charge. Ions are charged particles too as they have lost or gained electrons. The movements of these ions are responsible for the conduction of electricity. 

1.48 understand why covalent compounds do not conduct electricity

Covalent compounds do not have spare free electrons that can move and carry the charge; and neither do they contain ions-it’s a covalent compound not an ionic compound.

1.49 understand why ionic compounds conduct electricity only when molten or in solution

When they are a solid the ions are not free to move and carry the charge. When they are molten-as in it is in melted form-the ions are free to move. Remember in liquids the particles are able to slide over each other and move whereas in solids the particles can only vibrate around a fixed position. When the ionic compound is dissolved in a solvent to form a solution the ions are also made free to move.

1.50 describe simple experiments to distinguish between electrolytes and non-electrolytes

http://www.dummies.com/how-to/content/how-to-distinguish-electrolytes-from-nonelectrolyt.html

1.51 recall that electrolysis involves the formation of new substances when ionic compounds conduct electricity

Passing an electric current through a compound which is either molten or in solution causes chemical changes, the chemical reactions produce new products—new substances.

1.52 describe simple experiments for the electrolysis, using inert electrodes, of molten salts such as lead (II) bromide

Right hand electrode: Cathode-attracts cations-is the negative electrode, as it is attracting positive ions.

Left hand electrode: Anode-attracts anions-is the positive electrode, as it is attracting negative ions.

[Opposite charges attract]

Nothing happens until the lead (II) bromide melts.

Lead (II) bromide is an ionic compound. The solid consists of a giant structure of lead (II) ions and bromide ions packed regularly in a crystal lattice. It doesn’t have any mobile electrons, and the ions are locked tightly in the lattice and aren’t free to move. The solid lead (II) bromide doesn’t conduct electricity.

As soon as the solid melts, the ions do become free to move around, and it is this movement that enables the electrons to flow in the external circuit.

Electrodes are made out of carbon-which is inert (unreactive).

As soon as you connect the power source, it pumps any mobile electrons away from the left-hand electrode towards the right-hand one. The excess of electrons in the right-hand electrode makes it negatively charged-called the cathode. The left-hand electrode is positively charged because it is short of electrons. There is a limit to how many electrons can squeeze into the negative electrode (Cathode) because of the repulsion by the electrons already there.

The positive lead (II) ions are attracted to the cathode, which is the negative electrode. When they get there, they gain 2 electrons each from the electrode and forms neutral lead atoms. These fall to the bottom of the container as molten lead.

Pb²⁺ (l) + 2e^- à Pb (l)

This leaves spaces in the cathode that more electrons can move into. The power source pumps new electrons along the wire to fill those spaces.

Bromide ions are attracted to the positive anode. When they get there, the extra electron which makes the bromide ion negatively charged moves onto the anode because this electrode is short of electrons.

The loss of the extra electron turns each bromide ion into a bromine atom. These join in pairs (bond covalently) to form bromine molecules. Overall:

2Br^-(l) à Br₂ (g) + 2e^-

The new electrons on the anode are pumped away by the power source to help fill the spaces being created at the cathode.

The ions are discharged at the electrodes. Discharging an ion simply means that it loses its charge-either giving up electron(s) to the electrode or receiving electron(s) from it.

Redox reaction

Look back at the ionic equations, see that the lead (II) ions gain electrons at the cathode. Gain of electrons is reduction. The lead (II) ions are reduced to lead atoms.

The bromide ions lose electrons at the anode. Loss of electrons is oxidation. The bromide ions are oxidized to bromine molecules.

1.54 write ionic half-equations representing the reactions at the electrodes during electrolysis

For the electrolysis of lead (II) bromide, PbBr₂

Cathode:

Pb²⁺ (l) + 2e^- à Pb (l)

Anode:

2Br^-(l) à Br₂ (g) + 2e^-

Oxygen and Oxides

Note: This has just been updated as of 5/08, I expanded and changed bits on Spec 2.25, right at the bottom of this post. :) 


2.16 recall the gases present in air and their approximate percentage by volume

This is for unpolluted, dry air. Normally air contains a little water vapour too!

2.17 describe how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to determine the percentage by volume of oxygen in air

An experiment was carried out to find the percentage of air that is oxygen.  100 cm³ of air was passed from side to side over copper that was being heated with a Bunsen.  All the oxygen in the air will react with the copper.  No air could get in or out of the system while it was passed to and fro between the syringes.  As it was passed to and fro, the volume of air went down.  It was passed until the volume stopped decreasing, and a few minutes later the volume of remaining air was recorded.  There was 79 cm³ left. This shows that 21cm3 of the original 100cm3 of air was oxygen, because it was the oxygen that reacted with the copper to form black copper oxide. During this experiment, you should see the copper go black as it forms copper (II) oxide.

copper + oxygen à copper (II) oxide

2Cu (s) + O₂ (g) à 2CuO (s)

http://www.gcsescience.com/w10-air-oxygen.htm

2.18 describe the laboratory preparation of oxygen from hydrogen peroxide

Oxygen is most easily made in the lab from hydrogen peroxide solution using manganese (IV) oxide as a catalyst. The reaction is known as the catalytic decomposition (splitting up using a catalyst) of hydrogen peroxide. 

2H₂O₂ (aq) à 2H₂O(l) + O₂(g)

The oxygen produced is collected in an inverted glass cylinder by the downward displacement of water in a trough.

2.19 describe the reactions with oxygen in air of magnesium, carbon and sulphur, and the acid-base character of the oxides produced

When a substance burns in air, it reacts with oxygen gas and is said to be oxidised. 

2Mg (s) + O₂ (g) à 2MgO (s)

Magnesium burns with a bright white flame to give a white, powdery ash of magnesium oxide. The product is basic.

C (s) + O₂ (g) à CO₂ (g)

Carbon burns with a yellow flame, to give colourless carbon dioxide. The product is acidic.

S (s) + O₂ (g) à SO₂ (g)

Sulphur burns with a bright blue flame to give colourless sulphur dioxide, the product is acidic. Remember, sulphur dioxide gas is poisonous, and forms acid rain, so it's acidic! 

Any metal forms basic oxides, any non-metal forms acidic oxides. 

2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

The marble chips are basically calcium carbonate. Remember acids react with carbonates to give a salt, carbon dioxide and water. 

Acid + metal carbonate à salt + carbon dioxide + water

 CaCO₃ (s) + 2HCl (aq) à CaCl₂ (aq) + CO₂ (g) + H₂O (l)

It displaces the air as it is denser. 

2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper (II) carbonate

Most carbonates decompose (split up) when heated to from the metal oxide and carbon dioxide. 

In the case of copper (II) carbonate, this is a green powder which decomposes to give copper (II) oxide, which is a black powder and carbon dioxide when heated. 

CuCO₃ (s) à CuO (s) + CO₂ (g)

2.22 recall the properties of carbon dioxide, limited to its solubility and density

Carbon dioxide is a colourless, odourless gas, denser than air and slightly soluble in water. 

2.23 explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

It is used in carbonate (fizzy) drinks because it dissolves in water under pressure and when you open the bottle/can, the pressure falls and the gas bubbles out of the solution.

It is used in fire extinguishers because it is a dense gas, denser than oxygen and can displace it, basically pushing it out of the way so that no more oxygen reaches the fire. Thereby putting it out. 

2.24 recall the reactions of carbon dioxide and sulphur dioxide with water to produce acidic solutions

Rain is naturally slightly acidic because of dissolved carbon dioxide. Sulphur dioxide makes it even more acidic.

2.25 recall that sulphur dioxide and nitrogen oxides are pollutant gases which contribute to acid rain, and describe the problems caused by acid rain

Acid rain is formed when acidic air pollutants such as sulphur dioxide and nitrogen dioxide dissolve in rainwater. The sulphur dioxide and nitrogen oxides mainly come from power stations and factories burning fossil fuels, or from motor vehicles. The acid rain produces many problems. 


Sulphur dioxide dissolves in water in the atmosphere to form sulphurous acid (H₂SO₃). In the presence of oxygen in the air, this acid is slowly oxidised to sulphuric acid (H₂SO₄).

Oxides of nitrogen also contribute to acid rain. In the presence of oxygen and water, nitrogen dioxide is converted to nitric acid.

nitrogen dioxide + water + oxygen à acid rain

4NO₂ (g) + 2H₂O (l) + O₂ (g) à 4HNO₃ (aq)

The pH value of unpolluted rainwater is usually slightly below 7. This is because carbon dioxide in the air dissolves in rainwater to form carbonic acid, which is a weak acid.

CO₂ (g) + H₂O (l) à H₂CO₃ (aq)

However, acid rain is much more acidic than rain that only contains carbonic acid. Acid rain has a pH value of 4 or less. Which is slightly less acidic than vinegar, which is at around pH 3. 


So, what are the effects of acid rain?

• Acid rain reacts with metals and with carbonates in marble and limestone (calcium carbonate). When this happens, metal bridges and stone buildings are damaged, even statues if they're made of limestone. 
• Acid rain can reduce the pH of natural water bodies from between 6.5 and 8.5 to below 4, which will kill fish and other aquatic life. The water is then too acidic to support life.
• Acid rain also leaches important nutrients from the soil and destroys plants. Without these nutrients, plant growth is stunted. In some cases, acid rain dissolves aluminium hydroxide (Al(OH)₃) in the soil to produce Al³⁺  ions, which are toxic to plants. Forests throughout most of Central and Eastern Europe have been destroyed in this manner by acid rain. The plants are literally 'sick' and dying. 

acid rain eroding a statue