Topic 4: Reversible reactions and equilibrium

Brief intro

 

Changing the temp. of tubes containing nitrogen compounds:

 

• Heating-> nitrogen dioxide (NO₂) gas forms->Colour of solution changes to dark brown
• Cooling-> nitrogen tetroxide (N₂O₄) gas forms->Colour of solution gets paler

 

Water and ice are at a dynamic equilibrium because the rate of freezing = rate of melting. They are also reversible reactions.

 

Controlling the _PH of shampoo protects skin and eyes.

 

Reversible changes 4A

Molecular collisions

 

Irreversible change- products can’t form back into the original products

E.g. methane-> (combustion) ->Carbon dioxide + water = irreversible

 

Reversible- when products can react to produce the original reactants e.g.  Evaporation, melting

• Depends upon conditions; concentration, temperature, concentration

 

Example of RR

• Heating ice-> liquid
• Water cools to below 0 degrees Celsius
• Ice reforms

 

Forwards reaction= reactants make more products

Backwards reaction= products make more reactants

 

Temperature and concentration affects direction of change:

 

• When the concentration of REACTANTS is higher, it is more likely for the reactants to collide (as there are more of them), and produce lots of products. Thus the FORWARDS reaction is more likely.

 

• When the concentration of PRODUCTS is higher, it is more likely that the products will collide with each other and decompose to produce lots of the original reactants. Thus the BACKWARDS reaction is more likely.

 

• When heat is added to the reactants, they have an advantage and will thus produce products at a higher rate than the products are producing reactants. This is the forwards reaction.

 

• When heat is added to the products, the products have the advantage and will thus produce reactants at a higher rate than the reactants are producing the products at. This is the backwards reaction.

 

Equilibrium 4B

Water and ice are at equilibrium (balanced) at 0 degrees Celsius.  E.g. H₂O (l) = H₂O (s) [the equilibrium point is at any point between these two extremes]

The number of water molecules freezing is at the same rate as the number of water molecules melting.

 

When a reversible reaction takes place in a closed loop system (no reactants/ products lost), a state of equilibrium can be reached.

 

Reaching equilibrium-Iodine example:

 

• Iodine = soluble in potassium iodide
• The iodine is dissolved in hexane before it is dropped into the test tube with the potassium iodide (aq)
• 2 layers form: one layer of hexane + iodine, other layer with potassium iodide.
• The aqueous potassium iodide and organic solvent of hexane do not mix together
• Shake tube-> some of the iodine transfers to the potassium iodide, changing the colourless potassium iodide into a pale colour
• Shaking again-> more iodine transfers->the colour gets darker
• Continue shaking->the distribution of iodine is equal between the hexane and the potassium iodide-> equilibrium has been reached

 

Rules of equilibrium

 

• At equilibrium, the concentration of reactant and products doesn’t change; none of the particle are lost/ gained- only the distribution of particles changes.

 

• Equilibrium state can occur from the products or the reactants. E.g. if you put the iodine in the hexane, it will eventually lead to equilibrium of both solutions; likewise, if you place the iodine in the potassium iodide, it will also lead to equilibrium in both solutions.

 

Dynamic equilibrium

 

When the solution of iodine in hexane was shaken, the movement was in one direction- the iodine particles were all moving from the hexane to the potassium iodide in the forwards direction.

 

However, when the potassium iodide started to increase too much in iodine concentration, the iodine particles began to head in the backwards direction (back to the hexane).

 

The forwards reaction is originally faster, but it gradually starts to slow, and then the backwards reaction starts to take precedence.

 

The rate becomes equal when the backwards and forwards reaction continue BUT there is no overall change because each layer is gaining and losing iodine at the same rate. The reactions cancel each other out. This is known as dynamic equilibrium.

 

Ammonia and green chemistry 4C

Ammonia is used in fertilisers and explosives. There is a high demand for these stocks. This means that many industries have to ‘fix’ nitrogen into nitrogen compounds, in order to make the fertilisers and explosives.

 

Benefits of fertilisers

• Less land is needed to grow crops ( food is manmade)-> supports towns + cities

 

Costs of fertilisers

• Overuse of fertilisers->Nitrogen is washed into seas->algae is formed->damaged ecosystems

 

The Haber process

Industrial conditions required:

• Catalyst- Iron
• Pressure: x200
• Temperature: 450 degrees Celsius

 

• Atom economy= 100% (all atoms are used up, there are no by-products).

 

• Yield= 15% BUT   yield can be increased by recycling un-reacted hydrogen and nitrogen. This way, more ammonia is produced using the same amount of reactants.

 

The reaction is reversible, but that doesn’t mean that all the hydrogen and nitrogen will be converted to ammonia- the gases don’t stay long enough in the vessel to reach equilibrium.

 

Synthesis of ammonia

 

• Air and natural gas are obtained
• – Air is processed->nitrogen

– Natural gas is cracked->hydrogen (+ CO, CO₂ waste products)

3)   Hydrogen and nitrogen go into the reactor; 450 degrees Celsius, catalyst, x200 pressure

4)  – Gas is cooled in the condenser->ammonia liquid

– Un-reacted hydrogen and nitrogen is recycled

 

Le Chateliers’ Principle

 

• A+B->C+D = an equilibrium reaction

 

• When you stress one side of an equilibrium reaction, the reaction will favour the side to relieve that stress.
• 4A + 4B-> C+ D = reaction is stressed on the A+B side because there is a higher conc. of it

 

• Le Chatelier’s principle states that the opposing side (C +D) will relive that stress in order to maintain equilibrium

 

• So, C+ D will take some of the A and the B in order to balance the reaction:

 

2A + 2B-> 2C + 2D = an equilibrium has been reached

 

• The concentration remains the same, but the distribution has changed.
• It has taken some of the reactants (A+B) and made it into products (C+D).

Pressure

There are four molecules of gas on the reactants side (x 1 nitrogen, x3 hydrogen’s).There are two molecules of gas on the products side (x2 ammonia)

 

The reactant side has a higher pressure because there are more molecules, and thus less space to move.

The product side has lower pressure because there are fewer gas molecules, so more space to move.

 

Le Chatelier’s principle states that the product side must do something to counteract the high pressure in the reactant side. So, the product side starts producing more Ammonia by using up the nitrogen and hydrogen in the reactant side. This means that the numbers of hydrogen and nitrogen molecules in the reactant side starts to decrease. When the number of molecules decreases, the pressure also decreases, until it reaches a balanced pressure.

 

Temperature

 

The reaction of Nitrogen and Hydrogen is EXOTHERMIC.

• Favoured by higher pressure and lower temperatures

 

The decomposition of Ammonia is ENDOTHERMIC

• Favoured by lower pressures and higher temperatures

 

Le Chatelier’s principle states that if the temperature of the equilibrium reaction rises in the surroundings, it needs to be counteracted by something that takes in energy to lower the temperature.

 

1. Hydrogen + nitrogen reacts and releases energy and forms ammonia = exothermic reaction.
2. There is lots of heat energy present in the surroundings. Too much heat/ high temp.->explosions
3. This reaction is counteracted by the ammonia which takes in the energy released from the exothermic reaction.
4. Because it takes in the energy, this lowers the temperature of the surroundings, but not too much.
5. The ammonia uses the energy to break down to form hydrogen and nitrogen.

 

Conditions in the industry

 

The conditions for the process are a compromise that balances efficiency with safety + cost.

 

Higher pressure= higher yield of ammonia

• BUT high pressure plants are expensive to build, thus the pressure is set as high as feasible, at x200 pressures are used.

 

Low temperatures=higher yield of ammonia

• BUT lower temperature means a slow rate of reaction

 

Nitrogen fixation

 

Nitrogen fixing- the process of turning nitrogen from air into nitrogen compounds e.g. ammonia- used for pharmaceuticals, fertilisers

 

Nitrogen can be a limiting factor in some areas (e.g. third world)-> no crops

 

Nitrogen’s triple bond makes it very stable and thus it is very difficult to simply take nitrogen from air.

For this reason, we need to ‘fix’ nitrogen to meet demands.

There are biological + non-biological ways of fixing nitrogen. E.g. industry, combustion, agriculture.

 

Biological ways

 

Nitrogen-fixing bacteria can fix nitrogen at room temperatures and pressures.

The enzyme, and therefore catalyst, Nitrogenase converts nitrogen into ammonia. (Contains Fe/Mo/S)

 

Other nitrogen fixing bacteria are found in rood nodules, legume plants, clovers.

 

Non-biological ways

 

Industrial nitrogen fixation->increases amount of fixed nitrogen in the environment->the natural Nitrogen Cycle is destroyed.

 

• Environmental: nitrate flows to sea->algal bloom->eutrophication + damaged ecosystems
• Health: nitrogen pollutes water->humans drink polluted water->serious health issues

 

Restoring the balance in the nitrogen cycle is vital for sustainable development.

 

Production

 

Ammonia production is a clean process b/c all reactants are used, atom economy=100%, recycled reactants.

• Mainly produces CO₂ and nitrogen oxides, which are taken in by plants and levels are reduced.

 

Future feedstocks

 

Hydrogen is extracted from methane via steam forming.

Methane fossil fuel sources are non-renewable (will run out), thus alternative ways of producing hydrogen should be thought about.

• g. electrolysis of water= cheap + renewable electrical supply (solar panels)-> reduces GHG’s

 

Catalysts

 

More efficient catalysts are being discovered, which allow reactions to take place at lower temperatures and pressures->money + energy saved.

E.g. Ruthenium catalyst:

• Yield: 20%
• x40 pressure ->more expensive BUT benefits outweigh

 

Chemists study nitrogen fixing from nature e.g. study Nitrogenase structure-> make artificial ‘enzymes’ that mimic natural enzymes- this allows them find out more data to use artificial enzyme at room temperature and pressure = cheaper, efficient, less time consuming, less energy used.