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Reaction rates and reversible reactions

Introduction

In this lesson we will discuss reaction rates and reversible reactions. Chemical reactions proceed at different speeds. Some reactions take a very short time to occur.For example white phosphorus ignites spontaneously when exposed to oxygen. Other reactions take place very slowly for example the rusting of Iron.

Quantitative analysis shows that most reactions slow down as the reactants are used up. Rate of chemical reactions are important in our daily lives as they are to industrialists and chemical engineers. Industrialists and chemical engineers are interested in obtaining products rapidly and easily and as cheap as possible. Therefore to achieve this reactions can be accelerated so that they are econonmically viable..

In this lesson we will discuss the rates of a chemical reaction and factors that affect them. The time that a reaction takes place is of great importance to the chemist in an industry. If the rate of reaction can be increased more products can be manufactured in a given time and hence lower the cost . In cases where a certain reaction takes place at an explosive rate it is necessary for the chemist to be able to control the speed of that reaction. Hence it is essential for the chemist to know;
How various conditions can affect the speed of a reaction
How to speed up the reactions or slow down .
 

In this lesson we will discuss the rates of a chemical reaction and factors that affect them.

The time that a reaction takes place is of great importance to the chemist in an industry. If the rate of reaction can be increased more products can be manufactured in a given time and hence lower the cost . In cases where a certain reaction takes place at an explosive rate it is necessary for the chemist to be able to control the speed of that reaction. Hence it is essential for the chemist to know;

How various conditions can affect the speed of a reaction How to speed up the reactions or slow down .
 

Laboratory Examples of very Fast and moderately fast reactions. The following video clip shows a very fast reaction and a moderately fast reaction. Click to play and observe what happens carefully.



 

In Experiment 1 as soon as the two reagents come into contact with each other a white precipitate of Silver Chloride is formed. This reaction takes place very fast.

NaCl(aq) + AgNO3(aq) w NaNO3(aq) + AgCl (s)
 

Experiment 11
A yellow deposit slowly appears. This is due to deposits of Sulphur. It takes a longer time for the product to appear compared to experiment 1.
The rate of the reaction is the rate of change of an amount or concentrations of a particular reactant or product per unit time.

Different reactions have different reactions rates . Some are very fast eg burning of magnesium and burning of charcoal while others are slow eg rusting.
 

Collision theory state that for a chemical reaction to take place particles (Ions, molecules) must collide in favourable orientation and possess more than a certain minimum amount of energy .

This minimum amount of energy is referred to as activation energy.
The smaller the activation energy the faster the rate of reaction and the greater the activation energy the slower the rate of reaction.

 

Reactions can only start when the reactants energy is equal to the activation energy required to start the reaction. Collision theory and activation energy explain why even exothermic reactions do not occur at room temperature. Collisions must be energetic enough to surmount the activation energy barriers. An exothermic reaction is self sustaining when it has started. This is because energy evolved makes more reactants to have activation energy. During this type of reaction more energy is produced from the formation of new bonds than the energy needed to break the existing bonds and therefore there is surplus of energy.


 

For an endothermic reaction energy in form of heat or light has to be absorbed making it difficult for the reactants to have activation energy.
This deficiency in energy has to be made up from an external source.
 

The method chosen in the laboratory to measure the rate of the reactions must make use of a change that is taking place in a reasonable amount of time.
The change should be measurable in terms of ;
Volume in case gas is produced,
Change in mass of reactants and products,
Time taken for a given mass to disappear,
Time taken for a certain amount of precipitate to form.

By the end of the lesson you should be able to


i. State examples of simple reversible reactions
ii. Identify the reactants and products in forward and backward reaction

In this lesson we will study reversible reactions.

In the above chemical reactions the first positive sign means react with and the second on the product side means 'and'
The arrow shows the direction to which the reaction proceeds . Therefore the whole chemical change is represented as;
Fe (s) + 2HCl(aq) w FeCl 2 (aq) + H 2 (g)
In the second example the reactants are sulphur and Oxygen while the products is Sulphur (IV) oxide since Oxygen gas is limited not excess. The reaction should be presented as
S (s) + O2 (g) w SO2(g)
Reversible reaction is a chemical reaction in which the products of forward reaction can combine to give back the original reactants.
 

Iron reacts with water to form Iron (III) oxide and hydrogen gas.

In the reaction the double arrow means reversible so that rate of forward and backward reactions are the same.
Forward reaction has Iron and steam as the reactants and tri Iron tetra Oxide and hydrogen as products.
Backward reaction has tri Iron tetra oxide and hydrogen as reactants and Iron and steam as products.
 

Calcium carbonate decomposes to Calcium oxide and carbon (IV) oxide gas.

 

In the equation when the reaction takes place in a closed system initially it decomposes to form lime (Calcium oxide) and Carbon (IV) oxide. This is the forward reaction.Backward reaction is favoured by accumulation of Carbon (IV) oxide that reacts with lime to form limestone (Calcium Carbonate) which is given off and a white solid is left behind.

Hydrated Copper (II) crystals lose water of crystallization as shown in the following equation.

When blue crystals of Copper (II) Sulphate are heated in a test tube steam is given off and a white solid is left behind.
When the tube is cooled and a few drops of cold water added the blue colour reappears and the tube becomes hot.

The following experiment is carried out to investigate the chemical equilibrium in Chromate (VI) ion. Click to play the video and observe what happens carefully.

Potassium chromate (VI) solution is orange. When a few drops of dilute hydrochloric acid are added to the solution it turns yellow due to the formation of dichromate (VI) ions (Cr 2 O7 2-). This makes the concentration of hydrogen ions to increase as the acid is added.
An equilibrium is established with more Cr 2 O7 2- in solution than Cr O42- hence orange colour is dorminant.
On adding a few drops of 2M Sodium hydroxide the colour of the solution turns to orange as shown in the equation.

 

On adding 2M sodium hydroxide OH- ions react with H+ ions to form water. This reduces concentration of H+ ions in the solution shifting the equilibrium to the left.
H+ ions are formed to replace the H+ ions used to form water.
Altering the concentration of any one of the components of the equilibrium mixture disturbs equilibrium by affecting the rate of forward and backward reaction.
This makes the reaction proceed in one direction until a new equilibrium is established.
Note that the arrow showing reversible at equilibrium as shown
Which is different to that of the normal reversible reaction.

When the rate of forward reaction equals the rate of backward reaction a state of balance is reached.



At first D and C dominates to form F and G and then as F and G accumulates reverse reaction starts building up until a state of balance is reached.
At the state of balance the reaction is said to have reached a dynamic equilibrium because reactants and products are simply co existing but there is constant interchange from products to reactants and vice versa.
The following experiment explains chemical equilibrium as a state of balance.

The video clip below shows how to investigate change of colour of indicator in acid alkali media as an example of reversible reactions.Click to play the video and observe what happens carefully.

 

Phenolphthalein indicator is colourless in water and acid.
It is pink in alkali
In reversible reactions products can form reactants unlike many reactions that are one way.
In a closed system where no reactants or products escape both forward and backward reactions occur at the same time.
As time goes both forward and backward reactions proceed at the same rate. This is called chemical Equilibrium or dynamic equilibrium. The individual atoms and molecules reacting and products do not change in concentration.

You will now learn about the factors affecting rates of reaction.

In the preparation of oxygen gas from hydrogen peroxide Manganese (IV) Oxide is normally used as a catalyst to increase the rate of reaction.
In the manufacture of ammonia by Haber process the following optimum conditions are needed to realize a high yield.
- A mixture of nitrogen and hydrogen in the ratio of 1:3 by volume
- A Pressure of about 200 ' 500 atmospheres
- A temperature of about 4000 ' 5000c
- Iron catalyst
In this lesson we shall study how this factors like catalyst, temperature, pressure, concentration (for aqueous reactants) surface area (if reactants are solids) and light effect the rates or reaction.

The effect on rate of reaction by any one factor can be determined in the laboratory if the other factors are kept constant.

To investigate the effect of concentration on the rate of reaction. Click to play the video and observe what happens carefully.


 

The table below shows sample results obtained from this experiment.


 

The graph below represents the concentration of sodium thiosulphate used against time for the cross to disappear.

From the graph, the higher the concentration of sodium thiosulphate solution, the less time taken for the black ink cross to disappear. From the collision theory, increasing the concentration of reactants increases the frequency of collision between reacting particles.The greater the number of collisions, the higher the rate of reaction.

 

The following experiment is carried out to investigate the effect of temperature on the rate of reaction.Click to play the video and observe what happens carefully.

 

The following shows sample results that were obtained.

The graph below shows effect of temperature on reaction rates.

The curve shows that the time taken for the completion of the reaction decreases drastically with the increase in temperature.
The rate of reaction increases because the velocity of all reacting particles increase as the temperature increases. This increases the kinetic energy of the particles. .
Increase in kinetic energy provides the particles with the necessary activation energy required for the reaction to occur.
For this reason many reactions in industry are carried out at high temperature to get better yield.
 

The following video clip was carried out to investigate the effect of surface area on the rate of reaction. Click to play the video and observe what happens carefully.

Calcium carbonate reacts with dilute Hydrochloric acid to form calcium chloride, water and Carbon (IV) oxide gas.
CaCO3(s) + 2HCl (aq) w CaCI2(aq) + CO2(g) + H2O(l) The Carbon(IV)oxide gas escapes through the cotton wool to the atmosphere.The cotton wool stops the acid spray from escaping during the chemical reaction.
A graph of loss in mass (g) or mass of Carbon (IV) oxide (g) evolved against time is shown below.

Loss in mass is faster for powdered marble and slow few large marble chips.
This means that the rate of reaction is faster for marble powder and slow for large marble chips. The more finely divided the solid, the greater the surface area.

 

The following experiment was carried out to investigate the effect of a catalyst on the rate of reaction. Click to play the video and observe what happens carefully. (Courtesy of You Tube)

Before manganese(IV) oxide is placed into hydrogen peroxide, the solution gives off oxygen slowly, but not enough to re-light a glowing splint.
When manganese (IV) oxide is added, the reaction becomes vigorous and oxygen is given off within a very short time. The oxygen is now enough to relight a glowing splint. The mass of manganese(IV) oxide before and after the experiment is approximately the same. Manganese (IV) oxide acts as a catalyst in this reaction.
It increases the rate of a chemical reaction but remains chemically unchanged at the end of the reaction. The following equation shows the products formed.
2H2O2(aq) + MnO2(s) w 2H2O(l) + O2(g)
 

If 3g are used instead of 1g, the amount of oxygen of a given time would be more, this being because of the large surface crew which would improve its efficiency. However the volume of oxygen gas produced would be the same.
The graph of the volume of oxygen given out against time with and without a catalyst are shown below.

 

A catalyst provides an easier path by lowering the activation energy of the reactants allowing the reaction to proceed faster.The graph below shows the activation energy of the reaction with a catalyst and without a catalyst.

The table below shows some catalysts and the processes they are involved in;

Roll over the mouse on each of the catalyst

Study illustration shown below to show the effect of pressure on gas particles.

Pressure is significant in reactions involving gaseous reactants. When pressure increases, the volume in which the particles are contained is reduced. This brings the particles closer thus increasing collisions and affects the rate of reaction
When the pressure decreases the volume increases creating more space for the particles hence reducing the rate of collision and consequently the rate of reaction.

 




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