GCSE OCR Gateway Chemistry C5.3:a-c Equilibria and ammonia

C5.3 Equilibria

Summary

Common misconceptions

In a reaction, when the rate of the forward reaction equals the rate of the backwards reaction, the reaction in a closed system is said to be in equilibrium.

Learners often do not recognize that when a dynamic equilibrium is set up in a reaction the concentration of the reactants and products remain constant. They think that they are equal. Learners also sometimes perceive a dynamic equilibrium as two reactions.

Underlying knowledge and understanding

Learners will be familiar with representing chemical reactions using formulae and using equations.

C5.3a To recall that some reactions may be reversed by altering the reaction conditions

C5.3b To recall that dynamic equilibrium occurs in a closed system when the rates of forward and reverse reactions are equal

Higher tier:

C5.3c To be able to predict the effect of changing reaction conditions on equilibrium position and suggest appropriate conditions to produce as much of a particular product as possible

Le Chatelier’s principle concerning concentration, temperature and pressure

Ammonia and Equilibria

For many years, chemists have illustrated how changing conditions: temperature, concentration and pressure, affect the position of chemical equilibrium using ammonia production by the Haber Process. 

First, what is ammonia?

Ammonia is a gas at r.t.p.

It is very soluble in water: 1ml of water can dissolve 777ml of ammonia gas.  You can watch ammonia and the fountain experiment here.  

Bromothymol blue is the indicator, blue in alkali.

Ammona is toxic

It is less dense than air at 0.73kg/m³ (Air  =  1.225 kg/m³)

Ammonia is a weak alkali

NH₃      +     H₂O       ⇌      NH₄⁺       +    OH^—

As a weak alkali, it can neutralise common acids and form soluble salts that find use as water soluble fertilisers, as example would be ammonium nitrate: NH₄NO₃

NH₄OH      +     HNO₃       ⟶      NH₄NO₃        +    H₂O

As an alkali it is also used in low concentrations as a treatment for stings and bites e.g. from ants and/or nettles because the stinging agent is formic/methanoic acid HCOOH.

It is used as a “mordant” in hair dye: remember the corny British hair dye advert that used Penelope Cruz speaking the strapline “No ammonia!!”

Ammonia is formed from covalent molecules. 

The bonding is covalent.  The bonds between nitrogen and hydrogen are pairs of shared electrons. 

Its structure is simple molecular.  The discrete groups of atoms are loosely bonded to each other with very weak van der Waals forces.

The effect of forming ammonia is to fix nitrogen from its molecular form (N2 ) into an accessible molecular form that plants can absorb through their root systems. Most plants ( except legumes) cannot absorb molecular nitrogen from the air and absorb it through their roots in the form of ammonia and ammonium ions and nitrate ions.

Second, fixing nitrogen in ammonia.

Nitrogen is fixed into ammonia from nitrogen and hydrogen using the Haber process. 

The Haber process is an equilibrium reaction.

N₂(g)   +    3H₂(g)       ⇌       2NH₃ (g)

⇌ this sign indicates not just a reversible reaction but an equilibrium state.

In these reactions, the rate of the forward reaction equals the rate of the backwards reaction, but the reaction must be in a closed system for it to be said to be in equilibrium.

A closed system is a reaction vessel in which nothing can get in and nothing can get out, like a stoppered bottle in the diagram below. 

Look at a reagent bottle half full of water at room temperature.

Inside the sealed bottle there are two processes happening simultaneously.

Water molecules from the liquid state enter the gaseous state at the rate r1.

Simultaneously, water molecules from the gaseous water vapour enter the liquid state at the identical rate r2.

r1  =  r2

How do we know that r1 = r2?

The water level never changes at constant temperature.

If we left the bottle on the lab bench for 10,000 years the water level would remain the same at constant temperature (that’s unlikely by the way given the rate of enhanced global warming!!)

And yet the water level is constantly changing its composition as molecules leave and enter the liquid state!!!

The level looks static yet it is changing all the time!!.

The contents of the bottle look static yet they are in dynamic equilibrium.

The water molecules (billions upon billions of them it has to be said) are in constant motion entering and leaving the liquid state. 

The concentration of water in the bottle never changes of course.

This is what it is if the volume of water is 500ml.

500ml weighs 500g at r.t.p. because water’s density is 1g/cm^3.

The molar mass of water is 18 g/mol.

So the number of moles of water is 500/18  =  27.78 moles.

That's for half a litre of water so the concentration of water in moles per litre will be 27.78  × 2   =   55.6M.

The concentration of water  [H₂O]  =   55.6M and is constant at constant temperature.

But the concentration of the water vapour is very much lower than this.

The two species, water and water vapour, are not at the same concentration even though we are looking at an equilibrium system.

However, the two opposing rates of evaporation and condensation are equal to each other  r1   =  r2.

All these conditions are true only if the stopper remains on the bottle but take it out and there is no longer a sealed system and no longer a perfect dynamic equilibrium.

These are the exclusive features of all perfect dynamic equilibria.

• constant temperature (think about what would happen if the bottle above was stood in a beaker of boiling water)

• sealed system (no material leaves or enters the reaction)

• equal and opposite reaction rates

• unchanging constant concentrations of reactants and products (sometimes referred to as the position of equilibrium)

Many reactions are reversible but very few exist in a perfect equilibrium state.

Chemists often assume the perfect equilibrium state exists for a reaction when in reality a very imperfect one exists especially in industrial processes where material is constantly being added at the start and removed at the end of the process!!

Let’s look now at the effect of temperature and pressure on the ammonia equilibrium and how the optimum conditions are created for its production.

We will use a practical principle that helps us decide the effect of changing conditions. Henri Louis le Chatelier, a French chemical engineer, first explained the principle back in the 1880’s.  

Le Chatelier said: 

“When a stress is applied to a system in equilibrium, the position of equilibrium shifts is such a way so as to absorb the change.”

Here’s how it works out:

A:  The effect of temperature change on the ammonia equilibrium

N₂(g)   +    3H₂(g)       ⇌       2NH₃ (g)

The forward reaction in this equilibrium—making ammonia—is exothermic and heat is given out (92kJ/mole to be precise).

So if we make the equilibrium hotter raising its temperature, Le Chatelier tells us that the equilibrium will seek to resist this change. 

It will change so as to make things cooler by favouring the backwards reaction i.e. the decomposition of ammonia—not its formation.  This means that at higher temperatures there is a lower yield of ammonia. 

This effect is not what we expect since we all think that if we warm up a chemical reaction we will create more product but not so here in this reaction.

B:  The effect of change of pressure on the ammonia equilibrium

More molecules     N₂(g)   +    3H₂(g)       ⇌       2NH₃ (g)   less molecules

The pressure exerted by a gas is a result of collisions of molecules with the surface of the container they are in.  You can see from the equation that the forward reaction—making ammonia— leads to fewer molecules in the reaction vessel. 

Now if we raise the pressure on the equilibrium Le Chatelier tells us that it will resist this imposed change.  The equilibrium will move its position to lower the pressure and that must mean fewer molecules in the reaction vessel and that must means more ammonia is produced at higher pressures. 

Higher pressures bring about greater yields of ammonia as they favour the forward reaction.

We can see the effect of changing pressure and temperature on the yield of ammonia in the graph below: 

C:  Creating the optimum (best) conditions for the production of ammonia.

N₂(g)   +    3H₂(g)       ⇌       2NH₃ (g)

First, a reasonably safe and cost effective pressure is used.  Too high risks explosion and also means greater expense in building an industrial plant to withstand a very high (say 800 atmos) pressure.

Second, a temperature is used that is also cost effective but not so low as to make the reactions of both forward and backward too slow and unprofitable.  So we use 400^oC. 

Third, to ensure that the reaction goes fast enough at this optimum temperature a finely divided, high surface area iron (Fe) catalyst is used.  The catalyst does not increase or decrease the ammonia yield but it does ensure the yield is achieved in as short a time as possible.

You can see the optimum conditions pointed out on the graph above.