Research consultancy
SENIOR FOUR TERM I
REACTION RATES AND REVERSIBLE REACTIONS
GENERAL OBJECTIVE
By the end of this topic, the learners should be able to explain the effect of different factors on reaction rates.
Specific objectives
Learners should be able to:
Define rate of reaction
Describe some methods used to measure rates of reactions
Illustrate reaction rates graphically and explain the representation quantitatively
Using experimental data.
Recall simple reversible reactions.
Recognize reversible sign and explain how reversible reactions reach a state of “balance.”
REACTION RATES
There are several reactions which take place in everyday like situations and in laboratory and industrial settings. From observations some of these reactions proceed more quickly than others. For example when a solution of potassium iodide is added to a solution of lead(II) nitrate, a deep yellow precipitate of lead iodide takes place instantaneously, but the rusting of iron exposed to moist air may take a few days. Radioactive substances may take a few seconds to millions of years to decay. In this topic, we shall investigate the rates of some reactions and the factors that affect them.
The reaction rate indicates the velocity of chemical reactions. Rate of reaction is the amount of reactants consumed or the amount of product formed per unit time. This can be expressed as:
Rate of reaction = (change in amount (concentration of substance)/(time interval)
The equation defines the rate of reaction as the rate of change in concentration of a particular reactant or product. Chemical reactions proceed at different rates.
For example, when magnesium dissolves in acid.
Mg_((s))+ HCl_((aq))→MgCl_2(aq) +H_(2(g))
The rate can be calculated by measuring: –
The amount of magnesium or hydrochloric acid used up in a certain time, or
The amount of magnesium chloride or hydrogen formed in a certain (unit) time.
Put Mathematically,
Rate of reaction=(amount of magnesium or HCL consumed)/(time interval)
Or
Rate of reaction=(amount of magnesium chloride or H_2 formed)/(Per unit time)
Factors affecting the reaction rates
The concentration of reactants
The pressure of reactants for gaseous reactants.
Temperature
Surface area of reactants (particle size)
Catalyst
Light
Measuring rates of reaction
The rate of a chemical reaction can be measured by measuring the following variables: –
Rate of production of a gas.
Change in mass of reactants.
Change in intensity of colour.
Formation or disappearance of a precipitate.
Gaseous products
Many reactions give a gas as one product. The rate of such reactions can be found by one of two methods: –
Measure volume of a gas produced in a given time.
Allow the gas to escape from system. Measure reduction in mass of system after a given time. This works best for heavy gases like carbon dioxide.
Method 1
Volume of gas
Reactive metal + acid →salt solution+hydrogen gas
e.g. Magnesium+hydrochloric acid →Magnesium chloride+hydrogen
Mg_((s) )+HCl_((aq))→MgCl_(2(aq))+H_(2 (g))
The apparatus used is shown below
Clock
As the reaction proceeds, the gas pushes it way out into the syringe, the plunger is forced back. Readings of volume are taken every minute, using the scale marked or tabulated the side of the syringe. The results can be shown graphically as in the table below.
Time (min) (Volume (cm3)
0 0
1 12
2 20
3 26
4 30
5 33
6 35
7 36.5
8 37.5
9 38
10 38
The reaction is much faster at the start: 12cm3 are produced during the first minute, but only 3cm3 during the fifth minute (33-30cm3) = 3cm3
Notice that the curve is steepest at the start: after 9 minutes, it has gone completely flat. This shows that the reaction has gone to completion.
Start After 9 minutes
So the rate of reaction changes during the reaction.
How to find rate of reaction
Rate of reaction can be found at any given time. This is by measuring the slope of the curve at that time. You do this by drawing a tangent to the curve.
For example; what is the rate of production of hydrogen after two minute?
To find the answer, following the steps: –
Draw a line to the curve from the 2 min mark.
Carefully draw a tangent to the curve, at the point where the line touches it. The slope of the curve at time = 2mins is the same as the slope of gradient of this tangent.
Measure the gradient of the tangent. From Mathematics we know that: –
Gradient=(change in vertical distance)/(change in horizontal distance)=(change in volume)/(change in time)
These steps are shown on the graph below
Method 2
This method allows the gas to escape from the system. Measure the reduction in mass of the system after a given time. This method works best for reactions producing heavy gases, like carbondioxide.
A suitable reaction would be:
Metal carbonate+acid →salt solution+carbondioxidgas
A good example would be using
CaCO_(3(s)) 2HCl_((aq))→CaCl_(2(aq))+CO_(2(g))+H_2 O_((l))
The apparatus needed is shown below
Apparatus for measuring rate of reaction that produces a heavy gas.
Measuring decrease in mass during reaction of calcium carbonate with dilute acid.
Procedure
Set up the apparatus as shown in the future but without calcium carbonate (marble chips) at first.
Add the calcium carbonate (marble chips) Place the lid on the beaker. Read the mass of the beaker and its contents and start the stop clock.
Record the mass after 1 minute interval until the reaction is over.
Explain this with a balanced chemical reaction.
Tabulate and record your work in the table like the one below.
Time (min) 1 2 3 4 5 6 7 8 9 10
Mass (g)
Loss in mass (g)
Plot a graph with time on the horizontal axis and loss in mass on vertical axis. Draw a smooth curve through as many of the points as possible.
What can you tell from the slope of the graph at any point?
Why is the graph?
Steep at the start of the reaction
Less steep in the middle of the reaction?
Horizontal at the end of reaction? At what time does the reaction stop?
(a) Calculate the average rate of reaction.
(b) The rate of reaction at 2 minutes.
Describe how the reaction rate changes with time.
Discussion/Explanation
2. Because the lid is loose, the gas can bubble out from the beaker; the beaker therefore weights less.
The equation of reaction is;
CaCO_(3(s))+2HCl_((aq))+CO_2(g) +H_2 O_((l))
The decrease in mass is due to the carbondioxide gas escaping.
3. The graph slopes down because the mass of the system decreases.
1min
The slope of the graph at any time shows that the rate of reaction decreases with time.
The slope of the graph is steepest at the start of reaction because the reaction is faster at the start (A b)
The slope is less steeping the middle because the rate of reaction is decreasing as reactants become more and more consumed (BC).
The slope of the curve becomes horizontal because the reaction has come to completion i.e. the reactants are exhausted (CD).
Precipitated products
Reactions that occur in solution. This is precipitated as a cloudy suspension. This can be illustrated by the action of dilute acid on sodium this salphate solution. A fine suspension of sulpher is slowly formed.
For example a dilute hydrochloric acid reacts with sodium thiosulphate according to the equation.
Na_2 S_2 O_(3(aq))+2HCl_((aq))+ S_((s))+SO_2+H_2 O_l
Sulphur is precipitated.
The rate of reaction can be measured as below: –
Pour a known volume of sodium this sulphate say 40cm3, into a beaker and a known volume of dilute acid into another beaker.
Mark across with a blue or black ink or a piece of paper.
Put the paper under the thiosalphate beaker so that you can see the cross through the liquid.
Quickly pour the acid into the sodium thiosalphates and immediately start the stop clock.
As the cloudy suspension forms, i.e. blots out the cross from view. Stop the clock the second the cross completely disappears.
Cross on paper viewed
through solution
Dilute acid and sodium thio sulphate solution Clock
The faster the cross disappears, the faster the reaction is going, When the concentration of the acid is varied, the cross disappears faster with a more concentrated acid and takes longer for a more diluted acid.
The effects of concentration on reaction rates
The rate of reaction is proportional to the concentration of the reactants. When the concentration of the reactants is increased the rate of reaction also increases. This relationship can be represented graphically to 5000 give the following typically graphs.
From the graph, it can be seen that: –
The reaction in each case does not occur at the same rate throughout.
The gradient of the graphs are highest at the beginning. This is because the concentration of the reactants are higher and the reaction proceeds faster.
The gradient decreases as the reaction progresses i.e. the reaction slows down. This is because the reactants are being used up.
A point is reached; the graph becomes a horizontal line. This is because the reactants are used up and the reaction has come to completion.
The gradient of the graph at any point is equal to the rate of reaction at that particular point.
Examples
To investigate rate of sulphur deposition.
To investigate the effect of concentration on the rate of reaction between sodium thiosulphate and hydrochloric acid.
Requirements
180cm3 of BA¬1 containing 0.2M sodium thiosulphate.
60cm3 of BA2 containing 2M hydrochloric acid.
5 glass beakers of capacity 250cm3 conical flasks capacity 250cm3.
Measuring cylinders of capacity 100cm3.
Stop clock.
Filter paper.
Procedure
1. (a) Mark across the blue or black ink on the filter paper.
(b) Using a measuring cylinder, measure out 50cm3 of BA1¬into a 250cm3 beaker conical flask.
(c) Add 10cm3 of BA2 to BA1¬ in the conical flask and at the same time start the stop clock.
(d) Swire the beaker gently for the solutions to mix well and place the beaker or conical flask with its contents on the filter paper on the cross.
(e) Watch the cross through the solution from above the beaker. Stop the clock the second the cross disappearsonal note the time taken for the cross to disappear.
(a) Measure out 40cm3 of BA1 and add 10cm3 of distilled water into the conical flask.
(b) Add 10cm3 of BA2 and at the same time start of the stop clock. Follow the procedure in 1 above. Again note the time taken for the cross to disappear. Repeat the procedure above and record your results in the table below.
Number of flask Volume of BA1 (cm3) Volume of water added (cm3) Volume of BA2 (cm3) Time in seconds 1/T(5-1 )
150010
2401010
3302010
5203010
6104010
BA1 reacts with BA2 according to the equation.
〖Na〗_2 S_2 O_(3(aq))+2HCl_((aq))→2NaCl_((aq))+S_((s)) SO_(2(g))+H_2 O_((l))
〖Na〗_2 S_2 O_(3(aq))+2HCl_((aq))→2NaCl_((aq))+S_((s))+H_2 SO_(3(aq))
Questions
Why does the cross disappear in this experiment?
Write as ionic equation for the reaction
Plot a graph of:
Volume of BA1 against time.
Volume of BA1 against 1/time
How does the ion concentration of BA1 (thiosalphate affect the time for the cross to disappear)
Determine the rate of reaction of BA1 and BA2 and determine its units.
2. Temperature
The rate of reaction increases with increase in temperature. Generally the rate of a chemical reaction doubles for every 100C rise in temperature.
Experiment
To investigate the effect of temperature on the rate of reaction between sodium thiosulphate and dilute hydrochloric acid.
While the effect of concentration on rate of reaction is being studied, the hydrochloric acid/sodium thiosulphate experiment can be repeated, this time holding the concentrations constant but varying the temperature. This can be done by hearing the reactants in a water bath.
Requirements
0.2M sodium thiosulpphate.
2M conical flasks or beakers (250cm3)
Measuring cylinder (100cm3 capacity)
Measuring cylinder (10cm3 capacity)
Filter paper for any white sheet of papers.
Stop clock
Thermometer.
Procedure
Mark cross on a white sheet of paper or filter paper.
Place a beaker or conical flask on the paper over the cross.
Using a measure cylinder, measure out 50cm3 of the sodium thiosulphate and transfer on the conical flask on the filter paper.
Using a 10cm3 measuring cylinder, measure out 10cm3 of hydrochloric acid and pout into thiosulphate and immediately start the stop clock. Swire the mixture and place them on the cross. View the cross from above the mixture and note the time taken for the cross to disappear.
Measure out another 50cm3 of sodium thiosulphate into a second conical flask. Heat the sodium thiosulphate to 300C. immediately add 10cm3 of hydrochloric acid and simultaneously start the stop clock. Swire the mixture and put them on the cross. View and again note the time taken for the cross to disappear.
Repeat the procedure (a) to (e) above, heating the thiosulphate to 40, 50, 60 and 700C
Record your results in the table below: –
No. Temperature Tim in seconds 1/T(5-1)
1T =
230
340
450
560
670
Questions
Plot a graph of time against temperature
From the graph in (a) above, determine how long it would take for the cross to disappear if the solution of sodium thiosulphate was heated to 450C
Explain the shape of your graph.
Plot a graph of 1/time against temperature.
From the graph in (b) above, determine how long it would take the cross to disappear if sodium thiosulphate were heated to 600C.
Explain the shape of the graph.
Basing on the shape of the graph, what conclusion can you draw from the experiment?
Using the graph (a) Determine the rate of reaction between sodium thiosulphate and hydrochloric acid at a temperature of 400C.
From this experiment you can see that the rate of reaction increase as the temperature increases.
3. Effect of surface area on rate of reaction
The rate of reaction increases with increase in surface area. The effect of surface area can be determined by using calcium carbonate and an acid.
CaCO_(3(s))+2HCl_((aq))→CaCl_2+CO_(2(g))+H_2 O_((l))
The volume of carbondioxide produced is measured. In this experiment, a fixed mass of lumps (chips) of calcium carbonate (marble chips) is reacted with excess hydrocholoric acid in a conical flask fitted with a gas syringe. The volume of carbondiozide produced with time is measured. The experiment is repeated with the same mass of powdered marble chips (higher surface area) using the same concentration of hydrochloric acid.
Two sets of results are produced.
The results show that the reaction is faster with powdered marble. The powder has large surface area than the marble chips.
Note: The volume of carbondioxide (60cm3) is produced but with powdered marble chips, it is obtained within a shorter time.
4. Effect of catalyst on rate of reaction
A catalyst is a substance that alters the rates of chemical reactions but is itself not chemically changed or consumed at the end of the reaction.
The rate of reaction increases with the catalyst used.
The effect of catalyst on the rate of reaction can be decomposition of hydrogen peroxide. Hydrogen peroxide decomposes when heated to produce oxygen. The decomposition is faster when lead(II) oxide is added and even faster when manganese(IV) oxide is added catalysts generally increase the rate of reactions and the process is called catalysis.
Typically graphs with and without use of a catalysis can be shown below: –
Here the maximum volume of 800cm3 is obtained faster (in only 3 minutes) when catalyst is used and much longer without the catalyst (8 minutes)
Experiment
To investigate the effect of a catalyst of the decomposition of hydrogen peroxide
Dilute 11cm3 of volume 10 volume hydrogen peroxide to 100cm3 in a beaker. This is approximately 0.1M solution of the peroxide.
Set up the apparatus as shown in the diagram below but without manganese(IV) oxide at first. Set the volume of gas syringes at 0cm3.
Add the manganese(IV) oxide and start a stop clock immediately.
Read the volume of gas collected in each gas syringe at 7 minutes intervals until each is complete.
Identify the gas and write a balanced chemical equation. Which gas syringe contains more of the gas at any one time?
Give an explanation for this observation. Describe the appearance of the catalyst before and after the reaction.
Tabulate the results on the same axis. Plot graphs of volume of gas (cm3) on the vertical axis against time (minutes) on the horizontal axis for two different sets of readings.
Draw conclusion about reaction rate and
The presence of catalyst.
Different amounts of the same catalyst
The absence of a catalyst.
Repeat steps (2 – 6) but replace 1g and 3g of manganese(IV) oxide with 5cm3 and 15cm3 of dilute 1M hydrochloric acid.
Draw a conclusion about reaction rate and the presence of: –
Manganese(IV) oxide
Dilute HCl
Note: Hydrogen peroxide may decompose by the following reactions.
H_2 O_2(aq) ■(room temperature@( slow reaction ) ⃗ ) 2H_2 O_((l))+O_(2(g)) poor yield
H_2 O_2(aq) ■(room temperature catalyst@( fast reaction ) ⃗ ) 2H_2 O_((l))+O_(2(g)) good yield
H_2 O_2(aq) ■(room temperature,acid catalyst@( slow reaction ) ⃗ ) 2H_2 O_((l))+O_(2(g)) poor yield
Investigate the effect of manganese(IV) oxide on hydrogen peroxide
25cm3 of 0.1MH2O2 solution
25cm3 of 0.1MH2O2
Solution and 1g of MnO2
25cm3 of 0.1MH2O2 solution and 3g of MnO2.
Examples of catalytic processes
Process Catalyst
Haber process for manufacturer of ammonia Finely divided iron
Contact process for manufacturer of sulphuric acid Vanadium(V) oxide
Oxidation of ammonia to nitric acid Platinum
Hydrogenation of unsaturated oils to form fats in the manufacturer of margarine Hickee
Hearing potassium chlorate
2KClO_3→2KClO+O_2 manganese(IV) oxide
5. Effect of pressure on rate of reaction
Pressure has little or no effect on volumes of solids and liquids but has a big influence on the reactions in vowing gases. Increase in pressure in gaseous reactions means increase in the concentration of the reactants, and hence an increase in the rate of reaction between gases in the rate of reaction between gases. This increase in the rate is due to more frequent molecular collisions between gas molecules or atoms.
Effect of light on rate of reaction
Light, like heat is a form of energy, dust as some reaction can be speeded up by heat, so some can be speeded up by light.
A mixture of hydrogen and chlorine gases explode when exposed to bright light.
H_(2(g)) Cl_(2(g)) □(→┴( light ) ) 2HCl_((g))
A precipitate of silver bromide or silver chloride darkens when light is shone on it; this is the basis of photographic film.
The photosynthesis process goes faster in bright sunlight.
Reactions like those which are speeded up by light are called photochemical reactions. Photo chemical reactions are not very common in the laboratory setting.
The collision theory (theory of reaction rates)
From the previous discussions, we are now aware that the rates of chemical reactions are affected by: –
Increasing the concentration of reactants in solution.
Increasing the temperature.
Increasing the surface area of a solid.
Increasing pressure of a gas.
Using suitable catalyst.
The collision theory attempts to explain these observations.
The collision theory is derived from the kinetic theory which states that particles (atoms, molecules or ions) of a substance in solutions, liquids and gases are in constant from place to place.
The movement of any particular particle constantly changes in both speed and directions owing to collisions with other particles.
The collision theory states that before two or more substances can react, their particles must collide. Some collisions result into a chemical change i.e. chemical bonds in reactant particles break and new bonds are formed to make products. The colliding particles need a certain minimum amount of energy, called activation energy before their chemical bonds can break and new bonds form. The collision theory explains the facts affecting rate of reaction.
Concentration: An increase in concentration leads to an increase in rate of reaction. Higher concentration means a larger number of particles in a given volume. If there are more particles, there will be more collisions. More collisions means a large number of duceful collisions. Therefore the reaction rate will increase.
Temperature: An increase in temperature leads to an increase in reaction rate. Rise of temperature increases the speed of the particles and therefore their energy. There are more collisions per second, and more colliding particles have the necessary activation energy for the reaction. Usually arise of 100C doubles the rate of reaction.
Surface area: A greater surface area means a larger number of particles are exposed. Therefore there will be a larger number of collisions between particles.
Catalyst: The catalyst lowers the amount of activation energy and therefore increases the number of successful collisions per second and the rate of reaction increases.
NITROGEN AND ITS COMPOUNDS
General objective
By the end of this topic, the learners should be able to appreciate the importance of nitrogen in natural and industrial process.
Specific objectives
Learners should be able to: –
How nitrogen is prepared in the laboratory.
Outline the properties of nitrogen.
Explain how nitrogen is isolated from air.
State the uses of nitrogen.
Explain how ammonia is prepared in the laboratory.
Explain the difference in chemical reactions of ammonia gas and ammonia in aqueous solution.
Explain how ammonia is manufactured.
List the uses of ammonia.
Explain the preparation and manufacturer of nitric acid.
Explain the reaction of dilute and concentrated nitric acid.
Outline the uses of nitric acid.
State the methods of preparation of nitrates.
Name the products when different metal nitrates are heated.
Test for nitrates in the laboratory.
Nitrogen cycle.
Pollution effects of nitrogen and its compounds.
Occurrence
Nitrogen occurs naturally
as a salt pelete (sodium nitrates) in mineral NaNO3
in all living things as proteins and vitamins
as common compounds e.g. ammonium carbonate (NH4¬)2CO3
It occurs as a diatomic molecule N_2 or N≡N
Its atomic number is 7 and electronic arrangement 2:5
Its oxidation number is 3 (i.e. it gains or shares 3 electrons to attain the noble gas electronic structure or configurations.
It belongs to group V of the periodic table.
Methods of preparation
Laboratory preparation
(i) From air
Air contains 78% nitrogen 21% oxygen and 0.03% carbondioxide by volume.
When carbondioxide and oxygen are removed, what remains is largely nitrogen.
Removal of carbondioxide
Carbondioxide is removed by passing it through sodium hydroxide solution which absorbs carbondioxide.
Equation of reaction which takes place
2NaOH_((aq))+CO_(2(g))→Na_2 CO_(3(s))+H_2 O_((l))
Removal of oxygen
Oxygen is removed by passing the remaining gas heated over heated copper in a furnace.
Oxygen reacts with heated copper to form copper(II) oxide.
Equation of reaction which takes place
2CU_((s))+O_2(g) →2CUO_((s))
Preparation of nitrogen gas
Requirements
The apparatus is shown in the diagram below: –
Universal indicator paper or blue litmus paper.
Splint
Copper turnings
Concentrated sodium hydroxide solution (4 molar)
Water
Lime water
Magnesium ribbon
2. From a mixture of sodium nitrite and ammonium chloride
A mixture of sodium nitrate (NaNO3) and ammonium chloride (NH4Cl) are heated together. These react to give ammonium nitrite, which decompose to give nitrogen gas.
Equations of reactions
NH_4 Cl_((aq)) + NaNO_(2(aq)) → NH_4 Cl_(2(aq)) + NaCl_((aq))
Ammonium chloride Sodium nitrite Ammonium nitrite Sodium chloride
10g 13g
The ammonium nitrite decomposes as below to give nitrogen and water.
NH_4 Cl_(2(aq)) → H_2 O_((l)) + N_(2(g))
Ammonium nitrate Water Nitrogen
Procedure
Dissolve about 10g of ammonium chloride and 13g of sodium nitrite in some water in around bottomed flask. Heat the flask gently and collect it over water a few jars of the nitrogen liberated.
Preparation of nitrogen from sodium nitrate and ammonium chloride
Other reactions that yield nitrogen
i) Oxidation of ammonium by heated copper
(See action of ammonium on heated copper(II) oxide
2NH_(3(g))+3CuO_((s))→3Cu_((s))+3H_2 O_((g))
ii) Reduction of nitrogen(II) oxide by when passed over heated copper turnings
2NO_(2(g))+2Cu_((g))→2CuO_2+N_(2(g))
iii) By the action of chlorine on excess ammonia
The ammonia reacts with chlorine liberating nitrogen and forming hydrogen chloride.
3Cl_(2(g))+8NH_3→N_(2(g))+ 6NH_4 Cl_((s))
These methods are less effective than heating ammonium nitrate.
Properties of Nitrogen
Physical properties
Nitrogen is colourless, odourless to steles gas
Almost soluble in water for example 2cm3 of nitrogen dissolves in 100cm3 of water.
It does not burn not support burning.
It is neither acidic nor support burning.
Slightly less dense than air.
These properties show that there is no simple chemical test for nitrogen.
Test for nitrogen
It has no effect on litmus paper, burning splint, lime water or potassium permanganate i.e. it gives negative tests used for all gases thus there is no simple chemical test for nitrogen.
The teacher can design experiments to carry out the above tests with nitrogen prepared from ammonium chloride and sodium nitrite as shown above.
Industrial preparation of Nitrogen
Prepared on industrial scale by fractional distillation of air. Oxygen is obtained in the process. Air is first purified by removing dust particles from it. Carbondioxide is removed first and then water vapour.
The remaining air is cooled and compressed into a liquid. The liquidified air is allowed to evaporate. Nitrogen has a lower boiling point (77K) at a standard atmospheric pressure than oxygen (90K)
Chemical properties of nitrogen
Reactions of nitrogen and oxygen with metals
Oxygen is very reactive and reacts with most metals and non-metals forming their oxides. However, nitrogen reacts with the reactive metals of group I and group II, (Na, K and Ca and Mg) forming nitrites.
Reaction with magnesium: –
When burning magnesium burns in nitrogen, it does so to form magnesium nitrite.
Equation of reaction
Magnesium + nitrogen → magnesium nitrite
3Mg_((s)) + N_(2(g)) → Mg_2 N_(2(s))
White solid
When water is added to this compound, ammonia is given off which can be recognized by its characterized chocking smell.
Magnesium nitrate + Water → ammonia gas + magnesium hydroxide
Mg_3 N_(2(s) ) + 6H_2 O_((l)) → 2NH_(3(g)) + 3Mg(OH)_(2(s))
Reaction with hydrogen oxygen
If a mixture of oxygen and hydrogen are electrically sparked a mixture of oxides of nitrogen are formed.
Nitrogen + oxygen ⇌ 2NO_(2(g))
then
2NO_((g) ) + O_(2(g)) → 2NO_(2(g))
Nitrogen monoxide nitrogen dioxide
or (nitrogen(II) oxide nitrogen(IV)oxide
Reaction with hydrogen (under industrial conductions)
It reacts with hydrogen in the presence of iron catalyst to form ammonia (see Haber process for the manufacturer of ammonia)
Nitrogen + Hydrogen ⇌ ammonia
N_(2(g)) + 3H_(2(g)) ⇌ 2NH_(3(g))
The unreactive nature of nitrogen
Nitrogen has atomic number 7, and electronic configuration of 2:5
This can be represented as;
Or
It has a valency of 3 or can share three electrons to attain the noble gas electron structure.
Nitrogen exists a diatomic molecule joined by three covalent bonds as: –
Or (_x^x)N≡N_o^o
Each line – represents a pair of electrons and is called a bond (covalent bond) thus the three covalent bonds between the two nitrogen atoms are very strong making nitrogen unreactive.
Uses of nitrogen
The main use of nitrogen is in the maintenance of ammonia by the Haber process. Ammonia is an important chemical in the manufacturer of fertilizers.
Plants use nitrogen indirectly in the form of nitrogenous compound for their growth and survival.
Liquid nitrogen is used as a refrigerant. Bull’s semen for artificial insemination is stored under liquid nitrogen (boiling point 72K (-1960C).
Used to prevent fire outbreaks in empty fuel tankers. When the tankers are emptied of fuel, they are filled with nitrogen (Why use nitrogen?)
Ammonia NH3
Ammonia NH3 is a compound of nitrogen and hydrogen.
It is a covalent compound with the molecular mass 17.
Its molecular structure is
N ̈
H H H
Its molecular formula is NH3.
Methods of preparation of Ammonia
Laboratory preparation
By hearting a mixture of ammonia chloride once calcium hydroxide (or any other ammonium salt with an alkali). Sodium hydroxide and potassium hydroxide are not used because they are too reactive.
Equation for the reaction
Ca(OH)_2(s) + 2NH_4 Cl_((s) ) → 2NH_2(g) + H_2 O_((l) ) CaCl_(2(s))
The gas is prepared by downward displacement or upward delivery because it is less dense than air.
The flask is arranged in a slanting position so that the water which condenses on the colder parts of the apparatus will not run back into the hot flask and break it.
Calcium oxide is used as drying agent.
The usual drying agents; concentrated sulphuric acid and calcium chloride are not used because they both react with ammonia i.e.: –
Sulphuric acid + ammonia → Ammonia sulphate
H_2 SO_(4(l)) + +2NH_(3(g)) → (NH_4 )_2 SO_(4(s))
Calcium chloride ammonia → complex compound
CaCl_(2(s)) + 8NH_(3(g)) → CaCl_2.8NH_(3(s))
Preparation of ammonia
In the laboratory where the apparatus like he one shown above is not a viable, the set up below can be used.
Procedure
Grind up the solids, calcium hydroxide and ammonium chloride, with a pestle and mortar. Assemble the apparatus as shown above making sure the necks of the flask slopes downwards. Heat the flask gently and collect the jars of the gas as may be required.
Properties of Ammonia
Physical properties:
Ammonia is a colourless gas with a pungent, irritating smell.
Is less dense than air
Boils at -33.50C and freezes at -77.80C.
Extremely soluble in water. The most soluble of all gases.
The solubility of ammonia can be illustrated by the fountain experiment shown below.
Experiment
To investigate the solubility of ammonia gas in water
Apparatus and chemicals
Thick-walled round bottomed flak of capacity 1000cm3.
Ammonia generated from ammonium chloride and calcium hydroxide.
Teat pipette
Glass tube
Rubber stopper with 2 holes
Clump stand
Water trough
Water coloured with universal indicator.
Procedure
Fit the flask with a rubber stopper carrying a rube and a teat pipette with some water in as shown below.
Place the tube in water.
Pinch the teat pipette to force water into flask and watch what happens.
Water will push in the flask as a fountain and will continue until the flask is full of water as it was formerly full of ammonia.
The fountain experiment
Explanation
The water pinched into the flask from the teat pipette dissolves so much ammonia that there is a partial vacuum in the flask. Atmospheric pressure forces water rapidly up the tube and enters the flask as a fountain. The litmus turns blue.
This illustrates the high solubility of ammonia in water.
The ammonia dissolves in water to form aqueous ammonia.
NH_(4(g)) + H_2 O_((l)) ⇌ NH_4 OH_((aq))
This solution turns red litmus paper blue. It is the only alkaline gas.
The gas is poisonous + number from (i) upto somewhere
Reactions of ammonia (chemical properties of ammonia)
Turns red litmus paper blue – the only alkaline gas.
With hydrogen chloride (test for ammonia)
When a gas jar of hydrogen chloride is inverted over that of ammonia dense white fumes are seen. The dense white fumes are due to ammonium chloride formed.
The equation for the reaction
HCl_((g)) + NH_(3(g)) → NH_4 Cl_((s))
Experiment
To investigate the reaction of ammonia and hydrogen chloride
Required
Gas jars, concentrated ammonia solution, concentrated hydrochloric acid.
Procedure
Pour a little hydrochloric acid (about 1cm3/into one gas jar and concentrated ammonia into another.
Shake the gas jaw.
Bring the mouths of the gas jars into contact.
State what is observed.
Reaction with sulphuric acid
Sulphuric acid reacts with ammonia to form ammonium sulphate.
H_2 SO_(4(l) ) + 2NH_(3(g)) → (NH_4 )_2 SO_(4(g))
Solution in water
When dissolved in water, the resulting solution is called aqueous ammonia or ammonia solution. It is a weak alkali.
NH_(3(g)) + H_2 O_((l)) → NH_4 OH_((aq))
Ammonia solution or aqueous solution
In Solution, it ionizes partially as: –
NH_4 OH_((aq)) ⇌ NH_(4 (aq))^+ + OH_( (aq))^-
Ammonium ion Hydroxide ion
The above reaction shows the basic nature of ammonia gas. It is the only alkaline gas.
The equilibrium in the above reaction lies so much to the left.
Experiment
To test the alkaline nature of ammonia gas
Required
Concentrated ammonia in a reagent bottle. Blue and red litmus paper.
Water
Procedure
Wet the red litmus paper with water.
Bring the wet litmus paper at the mouth of the reagent bottle containing ammonia.
State what is observed
Repeat the above procedure with a lamp blue litmus paper and state what is observed
Precipitation of metal hydroxides
Precipitation of metal hydroxide that soluble in excess to from complex ions.
Reaction with copper(II) ions
When copper(II) sulphate solution is added to ammonia solution, an insoluble pale blue precipitate of copper(II) hydroxide is formed by double decomposition reaction.
CuSO_(4(aq)) + 2NH_4 OH_((aq)) → (NH_4 )_2 SO_4(aq) + Cu(OH_(2))
Copper(II) + Aqueous Ammonium Copper(II) hydroxide
sulphate ammonia sulphate (pale blue precipitate)
When excess ammonia is added, a further reaction takes place and a deep blue (royal blue) solution is obtained which contains the complex tetra amine copper(II) ions.
Cu(OH_2 )_((s) ) + 4NH_3 → [Cu(NH_3 )_4 ]_( (aq))^(2+) + 2OH_( (aq))^-
Deep blue solution (tetra amine copper(II) ion)
This reaction is used to use for copper(II) ions. It is a confirmatory use for Cu^(2+) ions.
Reactions with zinc ions
When ammonia solution is added drop wise to a solution containing zinc ions, acohite precipitation of zinc hydroxide is formed.
Zn_( (aq))^(2+) + 2NH_(3(g)) + 2H_2 O_((l)) → Zn(OH)_2(s) + 2NH_(4 (aq))^+
White precipitate
When ammonia solution is added in excess, the white precipitate becomes clean.
Zn(OH)_2(s) + 4NH_3(aq) → [Zn(NH_3 )_4 ]_( (aq))^(2+) +2OH_( (aq))^-
Tetra amine zink(II) ion (This is soluble and forms a colourless solution)
This reaction is used together with sodium hydroxide solution, in qualitative analysis to confirm the presence of zink ions. Zink hydroxide is soluble in both aqueous ammonia and sodium hydroxide solution.
Precipitation of metal hydroxides that are insoluble in excess ammonia solution
Reactions with lead(II) ions
Pb_( (aq))^(2+) + 2NH_(3(g)) + 2H_2 O_((l)) → Pb(OH)_2 + 2NH_(4 (aq))^+
Lead(II) hydroxide white precipitate insoluble in excess
Reaction with iron(II) ions
Fe_((aq))^(2+) + 2HH_(3(g)) → Fe(OH)_2(s) 2NH_(4(aq))^+
Dirty green precipitate insoluble in excess
Reaction with iron(III) ions
Fe_( (aq))^(3+) 3HH_3 + 3H_2 O_((l)) → Fe(OH)_(3 (aq) ) + 3NH_(4 (aq))^+
Brown precipitate insoluble in excess
Reaction with aluminium ions
Al_((aq))^(3+) + 3NH_3 + 3H_2 O_((l)) → Al(OH)_3(s) + 3NH_(4(aq))^+
White precipitate insoluble in excess
The above tests or reactions are used in qualitative analysis to identify these cations.
Oxidation of ammonia
Ammonia undergoes a number of oxidation reactions as shown by the following reactions
Reaction with oxygen (combustion)
Ammonia burns in acids supply of oxygen to produce nitrogen and water.
4NH_(3 (g)) + 3O_(2 (g)) → 2N_(2 (g)) + 6H_2 O_((g))
Combustion of ammonia (combustion of ammonia)
Catalyst combustion of ammonia
In the presence of platinum catalyst, ammonia is oxidized to nitrogen(II) oxide (nitric oxide) and water.
4NH_(3(g)) + 5O_(2(g)) □(→┴( pt cat ) ) 4NO_((g)) + 6H_2 O_((l))
The nitrogen(II) oxide formed can be oxidized to nitrogen(IV) oxide (nitrogendioxide) which can be seen as brown fumes.
2NO_((g)) + O_(2(g)) ⇌ 2NO_(2(g))
Nitrogendioxide (seen as brown fumes)
Later, the gases turn white because ammonium nitrate is formed.
4NO_(2 (g) ) + O_(2(g)) +2H_2 O_((l)) → 4HNO_(3 (g))
NH_(3 (g)) + HNO_3 → NH_4 NO_(3 (s))
Reaction with copper(II) oxide
Ammonia reacts with oxides of metals low in the activity series. The oxide has to be heated and dry ammonia passed over it.
3CuO_( (g)) + 2HH_(3 (g)) → 3Cu_( (s)) + 3H_2 O_( (l)) + N_(2 (g))
Thus ammonia acts as a reducing agent itself is oxidized to nitrogen.
Lead(II) oxide reacts in a similar way.
3PbO_((s)) + 2HH_(3 (g)) → 3Pb_((s)) + 3H_2 O_( (l)) + N_(2 (g))
Reduction of copper(II) oxide by ammonia
Catalyst oxidation of ammonia
Reaction with chlorine
Ammonia burns spontaneously in chlorine and is oxidized to nitrogen and hydrogen chloride.
Equation of reaction
2NH_(3 (g) ) + 3Cl_2 → N_(2 (g)) + 6HCl_( (g))
When chlorine is used up, the flame goes out and dense white fumes of ammonium chloride is formed by the reaction between excess ammonia and already formed by hydrogen chloride.
i.e. NH_(3 (g)) + HCl_( (g)) → NH_4 Cl_( (s))
Reaction between ammonia and chlorine
Preparation of aqueous ammonia
Set up the apparatus as above and heat the flask gently (see preparation of ammonia).
The rim of the funnel is inserviced just above the surface of water in the beaker. When the ammonia gas is passed in the trough, water rise up the funnel as it dissolves.
The level in the beaker falls and immediately air enters the funnel.
After a time the water in the beaker will have acquired the smell of ammonia gas which dissolved in it. This solution is known as aqueous ammonia or ammonia solution.
When ammonia dissolves in water, it ionizes as;
NH_(3 (aq)) + H_2 O_( (l)) ⇌ NH_(4 (aq))^+ + OH_((aq))^-
So the solution contains hydroxide ions, OH^-. This gives the solution its alkaline reaction towards litmus and many other reactions resembling those of caustic alkalis Na^+ OH^- and K^+ OH^-.
It reacts like these alkalis and will precipitate insoluble metallic acids when mixed with solutions of metals, e.g.
3KOH_( (aq)) + FeCl_(3 (aq)) → Fe(OH)_(3 (s) ) + 3KCl_( (aq))
3NH_4 OH_( (aq)) + FeCl_3 → Fe(OH)_( 3 (s) ) + 3NH_4 Cl_( (aq))
Zinc and copper hydroxide will dissolve in excess ammonia solution to gives solution of complex amines (see precipitation of metal hydroxides discussed behind).
It will neutralize acids to form ammonium salts which can be crystallized out and resemble or dinary metallic salts.
NaOH_( (aq)) + HCl_( (aq)) → NaCl_( (aq)) + H_2 O_((l))
NH_(3 (aq) ) + HCl(aq) → NH_4 Cl_( (aq))
2NaOH_( (aq)) + H_2 SO_(4 (aq)) → Na_2 SO_(4 (aq)) + H_2 O_((l))
2NH_(3 (aq)) + H_2 SO_4 (aq) → (NH_4 )_2 SO_(4 (aq))
Ammonium salts are electrovalent compounds containing the ammonium ion. NH_4^+ in combination with a corresponding acidic ion, e.g. CI^-, NO_3^-,SO_4^(2-),CO_3^(2-)
The ammonia molecule NH_3 can combine with a hydrogen ion (protons by donation of its lone pair of electrons. It is regarded as a base.
NH_(3 (aq)) + H_((aq))^+ ⇌ NH_(4 (aq))^+
Industrial preparation of ammonia (manufacturer of ammonia)
Ammonia is manufactured by the Haber process or synthesis in which dry ammonia and hydrogen gases are mixed and heated in the presence of finely divided iron as a catalyst.
N_(2 (g)) + 3H_(2 (g)) ■(Fe cat@⇌) 2NH_(3 (g) ) DH=-92KJmol-1
Flow diagram for the manufacturer of ammonia
Optimum conditions
1. The heating is done at a temperature of 4000C – 5000C (low temperature) since there is exothermic.
2. The pressure used is 200 – 5000 atmospheres since the reaction precedes a reduction in volume.
3. The catalyst used is finely divided iron promoted by aluminium oxide.
The ammonia formed can either be;
Liquified or
Dissolved in water
Raw materials
Nitrogen is obtained by fractional distillation of air.
Hydrogen – from hydrocarbons containing 5 to 9 carbon atoms e.g.
C_6 H_14 + 6H_2 O_((g)) ■(→@heat) 6CO_(2 (g)) + 13H_(2 (g))
Hexane
Uses of Ammonia
In the manufacture of nitric acid
In the manufacture of plastics like nylon
As a source of hydrogen when heated
2NH_(3 (g)) → 3H_(2(g) ) + N_(2 (g))
In the softening of water. It removes temporary hardness by precipitating calcium ion from calcium hydrogen carbonate.
Ca_((aq))^(2+) + 2HCO_(3 (aq))^- +2OH_((aq))^- → CaCO_(3 (s)) + 2H_2 O_( (l)) + CO_(3 (aq))^(2-)
In the manufacturer of fertilizers like ammonium sulphate (NH4)2SO4
ammonium nitrate NH4NO3
Used in the production of sodium carbonate in the solvary process.
Ammonium chloride is used in the manufacturer of dry cells.
Liquid ammonia is used in large scale refrigerating plants such as ships and warehouse.
It is used in the making of explosives
Tests for ammonia
Forms dense white fumes of ammonium chloride if brought into contact with fumes of concentrated hydrochloric acid.
NH_(3 (g)) + HCL_( (g)) → NH_4 Cl_((g) )
Ammonium chloride forms white fumes
Turns moist red litmus paper blue. It is the only alkaline gas.
Nitric acid
Laboratory preparation
By heating a mixture of potassium nitrate with concentrated sulphuric acid.
Equation of reaction
KNO_(3 (s)) + H_2 SO_(4 (l)) ⇌ KHSO_4 + HNO_(3 (g))
Potassium Sulphuric Potassium hydrogen Nitric acid
nitrate acid sulphate
This reaction is a general one. Any metallic nitrate when heated with concentrated sulphuric acid gives off nitric acid.
NaNO_(3 (g)) + H_2 SO_(4 (aq)) → NaHSO_4 + HNO_(3 (aq))
¬¬¬N.B: KNO_3 is better than NaNO3 because the latter is hygroscopic.
Industrial preparation of nitric acid
Nitric acid is formed by the oxidation of ammonia and three stages are in valued.
Stage 1
Catalyst oxidation of ammonia to nitrogen(II) oxide (nitrogen monoxide)
The amount obtained from the Haber process and air are passed over heated platinum catalyst at 7000C and a pressure of atmospheres. The ammonia is oxidized by oxygen to nitrogen monoxide (nitrogen II oxide).
Equation of reaction
Ammonia + oxygen → nitrogen(II) oxide + water
4NH_(3 (g)) + 5O_(2 (s)) → 4NO_((g)) + 6H_2 O_((g))
DH= -950Kjmol-
This reaction is highly exothermic.
Stage 2
Oxidation of nitrogen(II) oxide to nitrogen(IV) oxide.
Nitrogen(II) oxide (nitrogen monoxide) combines with more oxygen to form brown fumes of nitrogen(IV) oxide (nitrogendioxide)
Equation of reaction
Nitrogen(II) oxide + oxygen → nitrogen(IV) oxide
2NO(g) O2 2NO2(g)
Or 2 volumes 1 volume 2 volumes DH+ -114KJ mol-1
This reaction is favoured by high pressure because it results in the reduction in volume (3 volumes to 2 volumes).
The pressure chosen is 4 atmospheres
Stage 3
Reaction of nitrogen(IV) oxide formed above is then dissolved in water in the presence of excess oxygen, forming nitric acid.
Equation of reaction
Nitrogen(IV) oxide + water → nitric acid
4NO_(2(g)) + O_(2(g)) + 2H_2 O_((l)) → 4HNO_(3 (aq))
The rate of this reaction is favoured by high pressure.
The concentration of the acid is about 75%. More concentrated acid is made by mixing the 75% acid with concentrated sulphuric acid. The heat evolved during mixing vapourises nitric acid, which condenses to fuming acid.
Summary: Raw materials; ammonia, air, water conditions; temperatures 7000C, pressure 4 atmosphere, catalysis platinum.
Properties of nitric acid
Nitric acid can react as
An acid
An oxidizing agent
As an acid
Dilute nitric acid is a strong monobasic acid and is completely ionized as: –
HNO_(3 (aq)) + H_2 O_((l)) → H_3 O_((aq))^+ + NO_(3 (aq))^-
It has a sour taste and turns wet, blue litmus paper red.
Reaction with metals
Unlike other acids, it rarely gives hydrogen with metals.
It however, reacts with magnesium to produce hydrogen gas and metal nitrate. This is when the acid is very dilute
Magnesium + very dilute nitric acid → magnesium nitrate + hydrogen
Mg_((s)) + 2HNO_(3 (aq)) → Mg(NO_3 )_(2 (s) ) + H_(2 (g))
Reactions
4NH_(3 (g)) + 5O_(2 (g)) □(→┴( platinum cat ) ) 4NO_((g)) + 6H_2 O
2NO_((g)) + O_(2 (g)) □(→┴( platinum ) ) 2NO_(2 (g))
4NO_2 + O_2 + 2H_2 O → 4HNO_(3 (aq))
With oxides
Dilute nitric acid reacts with metal oxide forming corresponding nitrates and water.
Calcium oxide + nitric acid → calcium nitrate
CaO_((s)) + 2HNO_(3 (aq)) → Ca(NO_3 )_((aq) ) + H_2 O_((l))
With metallic carbonates
This reaction produces carbondioxide a characteristic of all carbonates and hydrogen carbonates e.g.
Lead carbonate + nitric acid → Lead nitrate carbonate + water
PbCO_(3 (s)) + 2HNO_(3 (aq)) → Pb(NO_3 )_2 + CO_(2 (g) ) + H_2 O_((l))
With alkalis
Dilute nitric acid neutralizes sodium or potassium hydroxides forming corresponding nitrate and water.
Nitric acid + sodium hydroxide → sodium nitrate + water
HNO_(3 (aq)) NaOH_( (aq)) → NaNO_(3 (aq)) H_2 O_((l))
HNO_(3 (aq)) KOH_( (aq)) → KNO_(3 (aq)) H_2 O_((l))
It also neutralizes ammonia solution forming ammonium nitrate
HNO_(3 (aq)) + NH_4 OH_((aq)) → NH_4 NO_(3(aq)) + H_2 O_((l))
Properties of nitric acid as an oxidizing agent
Concentration or dilute nitric acid behaves as an oxidizing agent. The reduction products of nitric acid vary. The common ones are nitrogen(II) oxide, nitrogen(IV) oxide and to a small extence nitrogen(I) oxide and ammonium nitrate may be formed.
Reaction with copper
Concentrated nitric acid
Oxidizes copper to copper(II) nitrate and it itself reduced to nitrogen(IV) oxide.
i.e. Cu_((s)) + 4HNO_(3 (aq)) → Cu(NO_3 )_(2 (aq) ) + 2H_2 O_((l) ) + NO_(2 (aq))
moderately dilute nitric acid reacts with copper to give nitrogen(II) oxides.
3Cu_((s)) + 8HNO_(3 (aq)) → 3Cu(NO_3 )_(2 (aq) ) + 2NO_((g) ) + 4H_2 O_((l))
Reaction with lead, zink and tin
Warm concentrated nitric acid oxidizes these metals to their nitrates.
e.g. Lead + nitric acid → lead nitrates nitrogen(IV) oxide + water
Pb_((s)) + 4HNO_(3 (aq)) → Pb(NO_3 )_(2 (aq) ) + 2NO_(2 (g) ) + 2H_2 O_((l))
4Zn_((s)) + 10HNO_(3 (aq)) → 4Zn(NO_3 )_(2 (aq) ) + NH_4 NO_(3 (aq))
Sn + 4HNO_(3 (aq)) → SnO_(2 (s)) + 4NO_(2 (g)) + 2H_2 O_((l))
(critically note these differences)
Aluminium and iron
These materials are made passive because of the formation of the oxide layer which stops further reaction.
Reaction with iron(II) sulphate
It oxidizes green solution of iron(II) sulphate to yellow solution of iron(III) sulphate in the presence of concentrated sulphuric acid.
Iron(II) sulphate + concentrated sulphuric acid + nitric acid
6FeSO_(4 (aq)) + 3H_2 SO_(4 (l)) + 2HNO_(3 (aq)) → 3Fe_2 (SO_4 )_3 + 4H_2 O_( (l) ) + 2NO_( (g))
The nitrogen monoxide formed is then oxidized to nitrogen dioxide seen as brown fumes.
2NO_( (g)) + O_((g)) → 2NO_(2 (g))
Reaction with non-metals
With carbon
Addition of concentrated sulphuric acid to carbon produces brown fumes of nitrogen dioxide and carbondioxide.
With sulphur
Concentrated sulphuric acid oxidizes sulphur to sulphuric acid and is itself reduced to brown fumes of nitrogen dioxide.
S_((s)) + 6HNO_(3 (aq)) → H_2 SO_(4 (aq)) +) 6NO_(2 (g))
Uses of nitric acid
The nitric acid produced on the industrial scale is often reacted with ammonia, one of its raw materials. The product is ammonium nitrate.
NH_(3 (aq)) + HNO_(3 (l)) → NH_4 NO_3
This is an important fertilizer.
Nitric acid is also used in the manufacturer of explosives, such as trinitrotoluene (TNT) and nitroglycerine clay sticks are soaked in nitroglycerine to give dynamic a powerful explosive.
In the manufacturer of drugs. A 1% alcohol in solution of nitroglycerine is a medium used for heart diseases.
In the manufacturer of dye staffs.
Nitrates
Preparation of nitrates
Nitrates of alkali metals like sodium nitrate, NaNO_3, potassium nitrate KNO_3 are prepared by neutralization between these alkalis and nitric acids.
e.g. sodium hydroxide + nitric acid → sodium nitrate + water
NaOH_((aq)) + HNO_(3 (aq)) → NaNO_(3 (aq)) + H_2 O_( (l))
Potassium hydroxide + nitric acid → potassium nitrate + water
KOH_( (aq)) + HNO_( (aq)) → KNO_(3 (aq)) + H_2 O_( (l))
Zinc nitrate, lead nitrate and copper nitrate are prepared by the action of nitric acid on the metal, its oxide, hydroxide or carbonate.
e.g. ZnO_( (s)) + 2HNO_(3 (aq)) → Zn(NO_3 )_(2 (aq) ) + 2H_2 O_((l))
Properties of nitrate
Physical properties
All metals are soluble in water
All are white crystalline solids
Chemical properties
When heated with concentrated sulphuric acid, nitric acid is formed.
e.g. sodium nitrate + sulphuric acid → sulphuric acid + sodium hydrogen sulphate + nitric acid
NaNO_(3 (g) ) + H_2 SO_(4 (l) ) → NaHS_4 O_((aq)) + HNO_(3 (l))
They are decomposed by heat and the product depends on the position of the metal in the electrochemical series.
Group 1 nitrates (sodium nitrate and potassium nitrates)
When heated, they first melt to colourless liquids and then slowly decompose to give a yellow nitrite and oxygen gas.
E.g. Sodium nitrate → sodium nitrite + oxygen
2NaNO_(3 (s)) → 2NaNO_(2 (s)) + O_( (g))
Potassium nitrate → potassium nitrite + oxygen
N.B: Sulphur and carbon burn vigorously and brilliantly in these nitrates
2NaNO_(3 (aq)) + 2S → 2NaNO_(2 (s)) + SO_(2 (g))
2KNO_(3 (s)) + 2S → 2KNO_(2 (s)) + SO_(2 (g))
2KNO_(3 (s)) + 2C → 2KNO_(2 (s)) + CO_(2 (g))
They use the oxygen liberated on decomposition of the nitrate to form the oxides.
Ammonium nitrate
It behaves differently when heated. It forms nitrogen(I) oxide
NH_4 NO_(3 (g)) → N_2 O_( (g)) + 2H_2 O_( (l))
Lead(II) nitrate [Pb(NO_3 )_2
When heated, a conspicuous cracking sound is produced. It then melts and a reddish brown gas nitrogen(IV) oxide and oxygen are formed.
2Pb(NO_3 )_(2 (s) ) → 2PbO_((s) ) + 4NO_(2 (g) ) + O_(2 (g))
Lead oxide is reddish-brown when hot and yellow when cold.
(The cracking sound is called descripitation)
Other nitrates decompose in a similar way to lead nitrate
Copper(II) nitrate → copper(II) oxide + nitrogen(IV) oxide + oxygen
2Cu(NO_3 )_(2 (s) ) → CuO_((g) ) + 4NO_(2 (g) ) + O_(2 (g))
Magnesium nitrate → magnesium oxide + nitrogen dioxide + oxygen
2Mg(NO_(3 (s) ) → 2MgO_( (s) ) + 4HNO_(2 (g) ) + O_(2 (g))
Zink nitrate → 2ZNO_((s)) + 4NO_((s)) + 4NO_(2 (g)) + O_(2 (g))
Note
These other nitrates are hydrated and do not produce the cracking sound when heated. They melt first and dissolve in water of crystallization forming a solution. When water is evaporated, decomposition starts.
The oxides of copper and zinc ware characteristic colours.
2Zn(NO_3 )_(2 (s) ) → 2ZnO_((s) ) + 2NO_2 + O_(2 (g))
Yellow when hot and white when cold
2Cu(NO_3 )_(2 (s) ) → 2CuO + 2NO_(2 (g) ) + O_(2 (g))
black when hot or cold
Nitrates of less reactive metals like mercury and silver
These decompose on hearting to nitrogen dioxide gas, oxygen and a metal.
E.g. Silver nitrate → silver + oxygen + nitrogen dioxide
2AgNO_(3 (aq)) → 2Ag_( (s)) + O_2 + 2NO_(2 (g))
Mercury nitrate → mercury + oxygen + nitrogen dioxide
2HgNO_(3 (s)) → 2Hg_( (s)) + O_(2 (g)) + 2NO_(2 (g))
Note:
The effect of heat on nitrates is used to detect the nitrate ion NO_3^-. This can be detected by the brown fumes of nitrogen(IV) oxide produced during heating.
Test for a nitrate ion, NO_3^-
Brown ring test for the nitrate ion, NO_3^-
To a solution of a nitrate, added freshly prepared solution of iron(II) sulphate followed by concentrated sulphuric acid slowly down the sides of a tilted test tube.
A brown ring is formed at the junction of the solution with sulphuric acid.
Tilt the test tube to 450
The brown ring test
Alternatively
Mix equal amounts (about 0.5g of each) of potassium nitrate and iron(II) sulphate and dissolve and then treat as above.
Some important nitrates and their uses
Sodium nitrate – as a fertilizer and also in making other chemicals like nitric acid and potassium nitrate.
Potassium nitrate – in the manufacture of gun powders.
Ammonium nitrate – manufacture of explosives
The nitrogen cycle
Nitrogen is important to plants and animals for production of proteins. Plants can not, however, absorb nitrogen directly but in a form of nitrates. Nitrogen has to be converted into forms that plants can use and this conversion is called nitrogen fixation.
Converting nitrogen in the air to nitrates in the soil (nitrogen fixation)
There are natural and artificial methods of fixing atmosphere nitrogen.
Natural ways of fixing atmosphere nitrogen
Certain bacteria living in colonies in the roots of leguminous plants e.g. beans and peas convert atmosphere nitrogen to nitrates.
Some other bacteria which live as free organisms in the soil are also able to fix atmosphere nitrogen to nitrates.
Lightening can cause nitrogen to combine with oxygen to form nitrogen oxide. When this dissolves in water, nitric acid is formed. In the soil, nitric acid is converted to nitrates.
Artificial methods of fixing atmosphere nitrogen
Because many farming lands contain inadequate amounts of nitrogen in the soil, because of man’s removal of plants for food, inadequate drainage and weathering, artificial means of restoring nitrogen have been used. This is usually done by adding fertilizers; either natural fertilizers, such as cowdung or farm yard manure or artificial fertilizers e.g. ammonium phosphate and ammonium nitrate. The ammonium fertilizers are manufactured from ammonia, which is manufactured by the Haber process using hydrogen and atmospheric nitrogen.
When plants absorb nitrates, they form proteins which are used for growth and repair of worn out tissues in plants and animals.
Action of nitrifying and denitrifying bacteria.
Certain bacteria in the soil breakdown urea and dead bodies of plants and animals and convert their problems into ammonia. Nitrifying bacteria oxidize this ammonia into nitrates.
Denitrifying bacteria oxidize ammonia and convert it to nitrogen which is returned to the atmosphere (air).
The way in which elemental nitrogen is continually circulate in nature can be summarized below. This cycle is known as the nitrogen cycle.
The nitrogen cycle
QUESTIONS: TYPICAL UNEB QUESTIONS
(a) (i) Name the raw materials used in the manufacture of ammonia
(ii) Write an equation leading to the formation of ammonia.
(b) Explain how formation of ammonia is affected by: –
(i) Pressure
(ii) Temperature
(c) State another factor that affects the formation of ammonia.
(d) Dry ammonia was passed over heated copper(I) oxide until there was no further change. State what was observed (Qn. 11, UNEB, 2000)
(a) (i) Draw a labeled diagram to show how a dry sample of ammonia can be prepared from ammonium chloride in the laboratory.
(ii) Write an equation leading to the formation of ammonia
(b) Dry ammonia gas was passed over heated lead(II) oxide.
(i) State what was observed.
(ii) Write equation for the reaction that took place. (Qn. 13, UNEB 2002)
(a) Write equation to show the reaction between copper(II) hydroxide and dilute nitric acid
(b) Aqueous ammonia solution was added drop wise to a sample of copper(II) hydroxide until ammonia was in excess.
(i) State what was observed.
(ii) Write the formula of the final product. (Qn. 7, UNEB, 2002)
(a) Write an equation leading to the formation of ammonia on a large scale
(b) State any two conditions for the reaction
(c) Write equation for the reaction between ammonia and copper(II) oxide (Qn. 4, UNEB, 2003)
(a) Describe how a dry sample of ammonia can be prepared in the laboratory (diagram not required)
(b) Name the reagent that can be used to test for ammonia and state what would be if ammonia is tested with the reagent.
(c) (i) Draw a labeled diagram of the set up of the apparatus that can show that ammonia can burn in oxygen.
(ii) Write an equation for the combustion of ammonia in oxygen
(d) Dry ammonia was passed over heated copper(II) oxide
(i) State what was observed
(ii) Write an equation for the reaction (Qn. 11, UNEB, 2005)
Although nitrogen is generally unreactive, it readily reacts with burning magnesium ribbon.
Give a reason why nitrogen is generally inert
Burning magnesium reacts with nitrogen
Give a reason for the reaction
Write an equation for the reaction
Water was added to the product in (b). Write equation for the reaction
(Qn. 1, UNEB, 2005)
(a) With the help of equations, outline how a dry sample of ammonia can be prepared in the laboratory starting from ammonia chloride. (Diagram not required)
(b) Draw a labeled diagram of the set up of the apparatus to show that ammonia is very soluble in water.
(c) Using equations where possible explain why dry ammonia is passed over heated lead(II) oxide, a colourless liquid is formed and a grey solid residue remains.
(d) Ammonium nitrate dissolves in water according to the following equations
NH_4 NO_(3 (g)) + H_2 O_((l)) → NH_4 OH + HNO_(3 (aq))
Explain using equations, why extensive use of ammonia nitrate as a fertilizer can make the soil acidic. (Qn. 14, UNEB, 2009)
(a) When dilute nitric acid was reacted with copper, a colourless gas, G which turned brown when exposed to air was evolved.
Name gas G
Write the equation leading to the formation of the brown gas.
(b) Write an equation for the reaction that would take place if the brown gas was dissolved in water.
(c) State what would be observed if concentrated sulphuric was heated with iron(II) sulphate solution. (Qn. 8, UNEB, 2009)
State what would be observed if dilute ammonia solution was added drop wise to aqueous solution of: –
Zink sulphate
Aluminium nitrate
(b) Write an ionic equation for the reaction in (a) (ii). (Qn. 9, UNEB, 2009)
(a) Describe how nitric acid can be manufactured using hydrogen and nitrogen as raw materials (illustrate your answer with equations.)
(b) Write equations to show the effect of heat on
(i) NH4NO3
(ii) Zn (NO3)2
(c) Potassium nitrate was heated with concentrated, sulphuric acid. Write equation for the reaction that took place.
(a) (i) Draw a labeled diagram of the set up of the apparatus that can be used to prepare a dry sample of ammonia in the laboratory.
(ii) Write equation for the formation of ammonia
(b) Write equation for the reaction between ammonia and
(i) hydrogen chloride
(ii) lead(II) oxide
(c) State what would be observed if ammonia solutions was added to a solution of copper(II) chloride drop wise until in excess.
(d) On heating a mixture of ammonium sulphate and aqueous potassium hydroxide, ammonia gas was produced according to the following equation.
(NH_4 )_2 SO_(4 (s) ) + 2KOH_((aq) ) → K_2 SO_(4 (aq) ) + 2H_2 O_((l) ) + 2NH_(3 (aq))
When Xg of ammonia sulphate was heated with excess potassium hydroxide until there was no further change 425.5cm3 of ammonia gas was evolved at s.t.p. Calculate X.
(S=32,K=39, H=1,N=14,O=16,1 mole of gas occupies 22.4dm^3 at s.t.p.)
(Qn. 12, UNEB, 2011)
When aqueous ammonia was added drop wise until in excess to a solution containing a cation X, a white precipitate was formed which dissolved to give a colourless solution.
Identify X
Write the formula of the cation in the colourless solution.
Write an ionic equation leading to the formation of the white precipitate.
Name one other metal that when treated with aqueous ammonia would form a precipitate solution in excess ammonia.
State what would be observed if the metal ion you have named in (d) was treated with aqueous ammonia until in excess. (Qn. 10, UNEB, 2012)
(a) Nitrogen can react with hydrogen in the presence of a catalyst which is finely divided to form ammonia in the Haber process.
State the source of nitrogen
Name the catalyst used in the reaction.
Explain why the catalyst is finely divided.
Write equation for the reaction leading to the formation of ammonia in the Haber process.
(b) Write equation to show that ammonia can
(i) act as a reducing agent
(ii) burn in oxygen
(c) Ammonia obtained by the Haber process can be converted to nitrogen(II) oxide.
(i) Write an equation leading to the conversion of ammonia to nitrogen(II) oxide
(ii) State the conditions for the reaction
(d) Write equation(s) to show how nitrogen(II) oxide can be converted to nitric acid. (Qn. 12, UNEB, 2013)
(a) Write an equation for the reaction between oxygen and
Ammonia in the presence of heated platinum
Nitrogen monoxide
(b) State how the product in (a) (ii) can be converted to nitric acid.
(c) Write an equation and state the conditions for the reaction between nitric acid and
(i) sulphur
(ii) lead(II) oxide
In each case, state what was observed and write an equation for the reaction that took place when sodium nitrate was heated strongly.
Alone
As a mixture with concentrated sulphuric acid (Qn. 14, UNEB, 2014)
(a) When a nitrate of a metal Y was heated strongly, brown fumes were observed together with a residue which was reddish-brown when hot and yellow when cold.
Identify Y
Write an equation for the reaction which took place
(b) The residue from (a) was heated with dilute nitric acid. Write an equation for the reaction which took place.
(c) To the produce in (b), dilute sodium hydroxide was added drop wise until there was no furhter change. State what was observed. (Qn. 7, UNEB, 2014)
(a) (i) Name two substances from which nitric acid can be prepared in the laboratory.
(ii) Write an equation for the reaction between the substances you have named in (a) (i).
(b) Write an equation between fuming nitric acid and copper (Qn. 4, UNEB, 2015)
SULPHUR AND ITS COMPOUNDS
General objectives
By the end of this topic, learners should be able to appreciate the use of sulphuric and its compounds.
Specific objective
Learners should be able to: –
Explain the extraction of sulphur from its ores.
State the properties of sulphur
State the allotropes of sulphur
Describe how monodinic and rhombaic sulphur are prepared in the laboratory.
Outline how sulhur reacts with oxygen, carbon, concentrated nitric and sulphuuric acids.
Sub-topic sulphurdioxide
Learners should be able to: –
Explain the preparation of sulphurdioxide
State the properties of sulphur dioxide
Outline the uses of sulphur dioxide
Explain the pollution effects of sulphur dioxide on the environment.
Sub-topic sulphuric acid
Learners should be able to: –
Describe the manufacturer of sulphuric acid
State the uses of sulphuric acid
Explain the difference in chemical action between dilute and concentrated sulphuric acid.
Test for the sulphates in the laboratory.
Hydrogen sulphide
Learners should be able to: –
Name the reagents used in the preparation
Hydrogen sulphide in the laboratory.
Identify the gas from its smell.
Explain the pollution effect of hydrogen sulphide on the environment.
Occurrence
Sulphur is found in a combined state as sulphides e.g. iron pyrites (FeS2)I copper pyrites (PbS)I ; Zinc blende (ZnS)g galena (PbS)I unhydrite (CasSO4)I and gypsum (CaSO4∙2H2O)
Also found as a free element in volcanic areas like USA, Italy and Japan.
Extraction of sulphur
It is extracted by the Frasch process. This process mainly involves the melting of sulphur and then forcing it out of the ground.
Three concentric pipes of 15cm, 8cm, and 2cm in diameter are sunk in the sulphur deposit. Supper heated water at about 1700C and a pressure of 10 to 15 atmosphere is forced down the space between the two outer pipes. This melts the sulphur (mpt 1130C).
The molten sulphuric is forced out through the middle (8cm) pipe.
Molten sulphur is led into large tanks or rats and solidifies into yellow solid of over 99.8Y. purify or into molds to form sulphur.
Extraction of sulphur by Frasch process
Allotropies of sulphur
Sulphur like carbon exhibits allotropy i.e. it exists in more than one physical state. The chemical differences between the different allotropes is always small.
Sulphur has two allotropes; rhombic and monoclinic sulphur
Below 95.50C sulphur crystals exist in a form of rhombic sulhur and above 95.50C, crystals exists as monoclinic sulphur i.e. rhombic sulphur ■(above 95.5^0 C@■(⇌@below 95.5)) monoclinic
The temperature 95.50C is referred to as transition temperature i.e. the temperature at which one allotrope of the element changes to another.
Rings of rhombic sulphur consists of S8 molecules shown as below
Crystals of rhombic sulphuric
Rings of monoclinic sulphur
Crystal of monoclinic sulphur
Rhombic sulphur is octahedausal in
Structure and melts at 1140C. when rhombic sulphur is heated strongly, it is converted to monoclinic sulphur at 960C (this is the transition temperature for rhombic to monoclinic sulphur)
Rhombic sulphur has a density of 2.08gcm-3
Monoclinic sulphur has pencil like or needle like structure. It has a density of 1.98g cm-3. Its melting point is 1190C.
Laboratory preparation of rhombic or actahedral sulphur (∝-sulphur)
Required
Test tube
Filter paper
Funnel
Beaker
Carbond sulphide
Sulphur flowers or roll sulphur
Method
Shake some powdered sulphur with carbondisulphide ofr sometimes in a test tube.
Filter the contents of the test tube into a dry beaker through the funnel. Fasten the filter paper over the mouth of the beaker.
Piece some holes in it. The carbondisulphide will slowly evaporates, depositing the crystals of sulphure. Examine their shape. This type of sulphur is called rhombic sulphur.
Preparation of monoclinic sulphur (β-sulphur)
Place some powdered sulphur in a large crucible and carefully melt by gentle heating.
Stir and add more sulphur and repeat until the crucible is full of molten sulphur. Allow the crucible to cool. Carefully piece through in two pieces thin solid crust that forms of the mass and immediately pour off the liquid sulphur inside. Remove the solid crust by cutting around the edge of the crucible with a knife. Observe the crystals that have formed on the underside of the crust.
Observe the crystals
They are needle-like crystals of monoclinic sulphur.
Uses of sulphur
Manufacture of sulphuric acid
Used as a tunpcide
Making calcium hydrogen sulphite Ca(HSO3)2 used to bleach wood in the manufacturer of paper.
Vulcanization of rubber – ordinary rubber is soft and sticky. It can be made hard and strong by hearing with sulphur at 1400C, a process known as vulcanization of rubber.
Manufacture of explosives and matches e.g. gun powder is a mixture of sulphur, carbon and potassium nitrate.
Manufacture of medicine e.g. sulphur tablets and many forms of drugs contain sulphur.
In the manufacturer of plastic flowers.
Manufacture of chemicals used in detergents.
Chemical properties of sulphur
It is a relatively reactive element and reacts with other elements.
Reaction with air (oxygen)
It burns readily in oxygen with a blue flame forming sulphur(IV) oxide (sulphur dioxide)
S_s + O_(2 (g)) → SO_(2 (g))
Some amount of sulphur(VI) oxide is also formed
2S_((s)) + 3O_(2 (g)) → 2SO_(3 (g))
Sulphur trioxide (sulphur(VI) oxide)
Reaction with metals
Forms metallic sulphides when heated with metals
e.g. Fe_((s)) + S_((s)) → FeS iron(II) sulphide
2Cu_((s)) + S_s → Cu_2 S copper sulphide
Zn_((s)) + S_((s)) → ZnS_((s) ) zinc sulphide
Reaction with concentrated sulphuric acid and nitric acids
It gets oxidized by the above concentrated acids to sulphur(IV) oxides and nitrogen(IV) oxide respectively.
S_((s)) + 2H_2 SO_(4 (l)) → 3SO_(2 (g)) + 2H_2 O_( (l))
S_((s)) + 6HNO_(3 (aq)) → H_2 SO_(4 (aq)) + 6NO_(2 (g)) + 2H_2 O_((l))
Note:
Sulphur is oxidized by sulphuric acid to sulphur dioxide and the acid is reduced to the same substance (SO_2).
Of the three molecule of sulphurdioxide, in the equation(i) above, one is the product of oxidation of sulphur and two are products of reduction of sulphuric acid.
Both sulphur dioxide and nitrogendioxide are poisonous. The experiment should be done in a fume cupboard.
Reaction with carbon
Sulphur and carbon do not react at ordinary temperatures but at high temperatures such as that found at an electric furnace. The reaction product is liquid carbondisulphuric. It is used as a solvent.
Equation of reaction
C_2 + 2S_((s)) → CS_2
Carbondisulphide
Reaction with hydrogen
When hydrogen is bubbled through molten sulphur, hydrogen sulphide is formed.
Sulphur dioxide: SO2
Laboratory preparation
Can be prepared in the laboratory by two methods.
From sodium sulphate and dilute sulphuric or hydrochloric acid
Sodium sulphite + dilute sulphuric acid → sodium salphate + sulphuric oxide + water
Na_2 SO_(3 (aq)) + H_2 SO_(4 (aq)) → Na_2 SO_(4 (aq)) + SO_(2 (g)) + H_2 O_( (g))
Or
Na_2 SO_(3 (aq)) + 2HCL_((aq)) → 2NaCl_((aq)) + SO_(2 (aq)) + H_2 O_( (l))
Ionically
SO_(3 (aq))^- + 2H_((aq))^+ → SO_2 + H_2 O_((i))
From copper turnings and concentrated sulphuric acid
Cu_((s)) + 2H_2 SO_(4 (l)) → CuSO_(4 (s)) + SO_(2 (g) ) + 2H_2 O_((g))
Laboratory preparation of sulphur(IV) oxide
From sodium sulphate
The gas produced is passed through concentrated sulphuric acid to dry it.
It is collected by downward deliverty.
When the gas jar is full of sulphur dioxide gas, the potassium dischromate paper turns from orange to green.
The same experiment set up above can be used to when concentrated sulphuric acid is poured on copper turnings. The mixture is heated and the gas evolved is dried and collected in the same way.
Properties of sulphur dioxide
Physical properties
Colourless gas with irritating chocking smell
Poisonous and should not be inhaled.
2½s times denser than air.
Readily liquefied
Highly soluble in water – fountain experiment can illustrate this property.
Turns most litmus paper red.
Chemical properties
Reaction with water
It dissolves in water to form weak sulphurous acid
H_2 O_((l)) + SO_(2 (g)) ⇌ H_2 SO_(3 (aq))
This acid turns blue litmus paper red.
It cannot be isolated because it decomposes as
H_2 SO_(3 (aq)) ⇌ H_2 O_( (l)) + SO_(2 (g))
In air, it can be oxidized to sulphuric acid
2H_2 SO_(3 (aq)) + O_(2 (g)) → 2H_2 SO_(4 (aq))
It forms a white precipitate with barium nitrate solution
Ba_((aq))^(2+) + SO_((aq))^(2-) → BaSO_(3 (s))
Barium sulphite
Note:
Barium sulphate is soluble in dilute nitric acid
Barium nitrate also forms a white precipitate with the sulphate ion (SO_4^(2-)) i.e. Ba_((aq))^(2+) + SO_(4 (aq))^(2-) → BaSO_(4 (s))
This does not dissolve in dilute nitric acid, thus the SO_3^(2-) and the SO_4^(2-) can be distinguished.
When chlorine is passed through a solution of sulhurous acid, it is converted to sulphuric acid.
2H_2 SO_(3 (aq)) + Cl_(2 (g)) → H_2 SO_(4 (aq)) + 2HCl_((aq)) + SO_(2 (g))
Reaction with oxygen
Reacts with oxygen in the presence of platinum or vanadium(V) oxide catalyze to form sulphurtrioxide (sulphur(VI) oxide)
2SO_(s (g)) + O_(2 (g)) ■(pt cat@⇌) 2SO_(3 (g))
White fumes
Reaction with alkalis to form hydrogen sulphites
With excess sodium hydroxide, sodium sulphate is formed.
2NaOH_((aq)) + SO_(2 (g)) → Na_2 SO_(3 (aq)) + H_2 O_((;))
Sodium sulphite
With excess gas, sodium hydrogen sulphite is formed
NaOH_((aq)) + SO_(2 (g)) → NaHSO_(3 (aq))
Excess
Test for sulphur dioxide
The gas can be tested by using acidified potassium dichromate(VI) solution which changes colour from yellow or orange to green.
Cr_2 O_(7 (aq))^(2-) + 3SO_(2 (g)) + 2H^+ → 2Cr_((aq))^(3+) + 3SO_4^(2-) + 2H_2 O_((l))
Yellow or orange green
It declourises acidified purple potassium manganate(VII) solution.
Oxidizing a cation
Reaction with mageniusm
Burning magnesium continues to burn in sulphur dioxide forming a white solid of magnesium oxide and sulphur.
2Mg_((s)) + SO_(2 (g)) → 2MgO_( (s)) + S_( (s))
Thus sulphur dioxide oxidizes magnesium to magnesium oxide and itself is reduced to sulphur.
Reaction with hydrogen sulphide
It oxidizes hydrogen sulphide to a yellow sulphur.
2H_2 S_((g)) + SO_(2 (g)) → 3S_((s)) + H_2 O_((l))
Bleaching action
When moist blue or red flowers are added to sulphur dioxide in a jar, the blue or red colours of the flowers are bleached (decolourized).
The bleaching action here is due to sulphurous acid formed between water found in flowers and sulphur dioxide.
H_2 O_( (l)) + SO_(2 (g)) ⇌ H_2 SO_(3 (aq))
The sulphurous acid is then oxidized to sulphuric acid by oxygen from the coloured material (dye) and hence bleaches it.
H_2 SO_(3 (aq)) + (dye + O_2 ) → H_2 SO_(4(aq)) + dye
Coloured Coloured
Uses of sulphur dioxide
In the contact process for the manufacture of sulphuric acid
As a bleaching agent for wood, silk, straw, sponges, paper and textiles.
Manufacture of calcium and sodium hydrogen sulphates.
Calculating hydrogen sulphate is used for the manufacture of paper and artificial silk.
The gas is poisonous and kills microorganisms and therefore used for fumigation of houses and clothing.
It destroys insects like white ants.
Used as a preservative for some liquids. It reacts with oxygen and prevents oxidation of liquids e.g. fruits and other food stuffs are therefore preserved.
Refrigeration – it is used as a coolant in refrigerators.
The pollution effect of sulphur dioxide on the environment.
Sulphuric acid
Manufacturer of sulphuric acid by the contact process
There are four stages involved
Stage 1
Burning of sulphur in air to produce sulphur dioxide.
S_((g)) + O_(2 (g)) → 2SO_(2 (g))
The sulphurdioxide is mixed with air and thoroughly cleaned (purified) to removal any dust or other chemicals e.g. Arservic(II) oxide that would poison the catalyst (platinum) and tender it less effective. Vanadium(V) oxide is mainly used because it is less susceptible to poison and less expensive.
Stage 2
The sulphur(IV) oxide obtained above is passed over heated catalyst (vanadium(V) oxide), V_2 O_3, heated at 450-500°C to produce sulphur trioxide (sulphur(VI) oxide).
2SO_(2 (g)) + O_(2 (g)) ■(□(→┴(〖 V〗_2 O_5 cat ) )@450-500°C)2SO_(3 (g)) DH=-ve
Stage 3
The sulphur(VI) oxide is then mixed with concentrated sulphuric acid in an absorption chamber to produce fuming sulphuric acid called “oleum.”
N.B:
Direct combination of sulphur(VI) oxide with water is not used because the mixture forms amist which is difficult to condense.
Reason: The reaction is highly exothermic.
H_2 SO_(4 (l)) + SO_(3 (g)) → H_2 S_2 O_(7 (l))
(oleum)
The oleum is then mixed with water to produce the ordinary sulphuric acid.
H_2 S_2 O_(7 (l)) + H_2 O_( (l)) → 〖2H〗_2 SO_(4 (l))
Oleum ordinary sulphuric acid
The contact process is summarized in the flow chart below
Properties of concentrated sulphuric acid
Physical properties
Colourless odourless, oily liquid
Has a density of 1.8 that of water
Soluble in water to give a considerable amount of heat when a solution is formed.
It is hydrogscopic (its use as a drying agent depends on this property)
Decomposes at 333°C (its boiling point)
H_2 SO_(4 (l)) → 〖SO〗_(3(g)) O_(2 (g))
Chemical properties
(i) Reacts with salts of weak acids to form sulphates. This action can be violent.
(ii) Does not change litmus or any other indicator to show acidic properties.
(iii) Displaces more volatile acids from their salts.
E.g. Hydrogen chloride from chlorides and nitric acid from nitrates.
Oxidizing properties
(i) Hot concentrated sulphuric acid is a strong oxidizing agnet. This accepts electrons or supplies oxygen to other compounds.
E.g. 2H_2 SO_4 → 2SO_(2 (g)) + 2H_2 O_((g)) + O_2 (used for oxidizing)
Or
2H_2 SO_4 + 2e → SO_(4 (aq))^(2-) + 2H_2 O_( (l)) + SO_(2 (g))
N.B: The electrons are supplied by the reducing agent concerned in the reaction. This may be a metal such as zinc or copper.
E.g. Cu (or Zn) → Cu^(2+) (or (Cu^(2+) ) + 2e
Then 2H_2 SO_(4 (aq)) + 2e → SO_(4 (aq))^(2-) + 2H_2 O_( (l)) + SO_(2 (g))
The metallic ion associates with the sulphate ion to form the corresponding metal sulphate.
i.e. Cu_((aq))^(2+) + SO_(4 (aq))^(2-) → CuSO_(4 (aq))
Overall equation
Cu_((s)) + 2H_2 SO_(4 (l)) → CuSO_(4 (s)) + 2H_2 O_( (l)) + SO_(2 (g))
Note:
This reaction is used to prepare sulphur dioxide in the laboratory.
Metals do not liberate hydrogen from concentrated sulphuric acid.
The sulphuric dioxide can be detected by decourization of potassium manganate(VII) solution.
(ii) Hot concentrated sulphuric acid can also oxidize can also oxidize non-metals like sulphur and carbon to their oxides.
E.g. S_((s) ) + H_2 SO_4 → 2H_2 O_( (l)) + 3SO_(3 (l))
C_2 + H_2 SO_4 → CO_(2 (g)) + SO_(2 (g)) + 2H_2 O_( (l))
Note:
The sulphur dioxide given off can be detected by using a filter paper moistened with potassium dichromate which turns from yellow/orange to green.
If hydrogen sulphide is bubbled through concentrated sulphuric acid, it is oxidized to sulphur.
H_2 S_( (g)) + H_2 SO_(4 (l)) → S_((s)) + 2H_2 O_((l)) + SO_(2 (g))
3. concentrated sulphuric acid as a dehydrating agent
It has a great affinity for water and they mix with a great evolution of heat. Equal volumes when mixed can produce heat of about 1200C. This shows a chemical reaction and heat is due to the hydration of ions which result from sulphuric acid.
It is hygroscopic (absorbs water from the atmospheres. Thus it is used as a drying agent.
Caution
When mixing, the acid with water, it is important not to add water into the acid but slowly add the acid to water.
Concentrated sulphuric acid is used to dry gases like SO_2,Cl_2,HCl, but it cannot be used to dry reducing agents like hydrogen sulphide (H_2 S) or alkali gas ammonia (NH_3).
Because concentrated sulphuric acid remove water from such substances, it is regarded as a drying agent.
Below are some reactions in which it acts as a dehydrating agent.
With sugar (sucrose)
When concentrated sulphuric acid id added to wet crystals of sucrose in a small beaker, the mixture becomes hot and steam is given off. A black spongy mass of carbon swells and may fill the beaker.
Beginning
After one minute
Reaction taking place
C_12 H_22 O_(11 (s)) + nH_2 SO_(4 (l)) → 12C + 11H_2 O_( (l)) + 2H_2 SO_(4 (l))
sucrose black mass of carbon
The carbon is then oxidized by sulphuric acid
C + 2H_2 SO_(4 (l)) → CO_(2 (g)) + 2SO_(2(g)) + 2H_2 O_( (l))
N.B: The corrosive nature of the acid is explained in the same manner.
Cotton is simply cellulose, whose formula is; (C_6 H_10 O_5 )_n
It reacts with sulphuric acid as;
C_5 H_10 O_5 )_((s)) + nH_2 SO_(4 (l)) → 6C + 5H_2 O + nH_2 SO_4
With copper sulphate: CuSO_4.5H_2 O
When blue crystals of copper(II) sulphate are warmed with concentrated sulphuric acid, the blue crystals become white.
CuSO_4∙5H_2 O_((s)) + H_2 SO_(4 (l)) → CuSO_(4 ( )) + 5H_2 O_((g))
Blue
Thus sulphuric acid removes the molecules of water of crystallization from hydrated copper(II) sulphate (blue) to the white unhydrous copper(II) sulphate.
With ethanol
Heating ethanol with sulphuric acid liberates ethane (an alkene).
CH_3 CH_2 OH + ■(□(→┴(Conc H_2 SO_4 ) )@170°C) CH_2=CH_(2 (g)) + H_2 O_( (g))
ethanol ethene
Acidic properties of dilute sulphuric acid
Has a sour taste and turns blue litmus red
Forms carbondioxide with carbonates and hydrogen carbonates.
E.g. Na_2 CO_(3 (aq)) + H_2 SO_(4 (aq)) → Na_2 SO_(4 (aq)) + CO_(2 (g)) + H_2 O_( (l))
2NaHCO_(3 (g) ) + H_2 SO_(4 (aq) ) → Na_2 SO_(4 (aq)) + CO_(2 (g)) + H_2 O_( (l))
CaSO_(4 (s)) + CO_(2 (g)) + H_2 O_( (l)) → CaSO_(4 (s)) + CO_(2 (g)) + H_2 O_( (l))
Note: The calcium sulphate formed in this reaction is insoluble (forms a solid). The solid formed forms a coating on the remaining carbonate and the reaction go on completion.
Reaction with alkalis
It reacts with alkalis to form salt and water only. It is adibasic acid and forms two types of salts;
Sulphates and hydrogen sulphates.
2NaOH_( (aq)) + H_2 SO_(4 (aq)) → Na_2 SO_(4 (aq)) + H_2 O_( (l))
(Normal salt)
NaOH_( (aq)) + H_2 SO_(4 (aq)) → NaHSO_(4 (aq))
(Acidic salt)
The produce formed above depends on the quantity of the reactants used.
With excess sodium hydroxide, sodium sulphate, a normal salt is formed (equation i)
With excess sulphuric acid an acidic salt, sodium hydrogen, sulphate (equation (ii) above) is formed.
Reaction with metals
Dilute sulphuric acid reacts with many metals liberating hydrogen gas (those above hydrogen in the electro chemical series).
E.g. Zn_((s)) + H_2 SO_(4 (aq)) → ZnSO_(4 (aq)) + H_(2 (g))
Mg_( (s)) + H_2 SO_(4 (aq)) → MgSO_(4 (aq)) + H_(2 (g))
Reaction with metal oxides
It reacts with metal oxides to form a salt and water.
ZnO_( (s)) + H_2 SO_(4 (aq)) → ZnSO_(4 (aq)) + H_2 O_( (l))
Uses of sulphuric acid
Below are some uses of sulphuric acid in order of their importance:
Manufacture of fertilizers e.g. ammonium sulphate, obtained by neutralization of sulphuric acid with ammonia.
2NH_(3 (aq)) + H_2 SO_(4 (aq)) → (NH_4 )_2 SO_(4 (aq))
Calcium hydrogen phosphate (super phosphate) obtained by direct combination of calcium phosphate and concentrated sulphuric acid.
Ca(PO_4 )_2(s) + H_2 SO_(4 (l) ) → Ca(HPO_4 )_(2 (s) ) + CaSO_(2 (s))
Manufacture of dye stuffs, paints, plastics and explosives.
Manufacture of artificial silk (rayon) and nylon.
Production of other chemicals such as metallic sulphates, hydrochloric acid, hydrofluoric acid and nitric acid.
Used in acid-lead accumulators (batteries)
In the manufacture of detergents and soaps.
In the extraction of metals including “picking” to clean metallic surfaces, to remove rust before galvanizing or tinning.
Sulphates
All sulphates are soluble in water except: calcium sulphate CaSO_4), Barium sulphate BaSO_4), lead sulphate PbSO_4) and silver sulphate (Ag_2 SO_4)
Preparation of sulphates
All soluble sulphates are prepared by reacting dilute sulphuric acid with metals, metal oxides or carbonates.
By neutralization of dilute sulphuric acid with an alkali or carbonate
E.g. Sodium sulphate
2NaOH_( (aq)) + H_2 SO_(4 (aq)) → Na_s SO_(4 (aq)) + H_2 O_((l))
Na_2 CO_3 + H_2 SO_(4 (aq)) → NaSO_4 + CO_(2 (g)) + H_2 O_( (l))
By action of concentrated or dilute sulphuric acid on a metal or metal oxide.
E.g. zinc sulphate
Zn_((s)) + H_2 SO_(4 (aq)) → ZnSO_(4 (aq)) + H_2 O_( (l))
ZnO_( (s)) + H_2 SO_((4)aq)) → ZnSO_(4 (aq)) + H_2 O_( (l))
All other soluble carbonates can be obtained in a similar way.
Insoluble sulphates
These are sulphates of calcium, lead, and barium. They are prepared by double decomposition or precipitation method by reaction between solutions containing the ions of the metal and sulphate ions.
E.g. Pb(NO_3 )_(2 (aq) ) + Na_2 SO_(4 (aq) ) → PbSO_(4 (aq) ) → PbSO_(4 (s) ) + 2NaNO_(2 (aq))
Note:
Lead nitrate in solution ionizes as
Pb(NO_3 )_( 2 (aq) ) → Pb_( (aq))^(2+) + 2NO_(3 (aq))^-
Sodium sulphate ionizes as;
Na_2 SO_(4 (aq)) → 2Na_((aq))^+ + SO_(4 (aq))^(2-)
On mixing the two solutions:
Pb_((aq))^(2+) + SO_(4 (aq))^(2-) → PbSO_(4 (s))
Barium ions and calcium ions react in a similar way.
Test for sulphuric acid and soluble sulphates
Both dilute sulphuric acid and soluble sulphates ionize in water to furnish sulphate ions.
i.e. H_2 SO_(4 (aq)) → 2H_((aq))^+ + SO_(4 (aq))^(2+)
Na_2 SO_(4 (aq)) → 2Na_((a))^+ + SO_(4 (aq))^(2-)
To test for the sulphate ion, we can use either barium chloride solution or barium nitrate solution. If barium chloride is used, it should be followed by dilute hydrochloric acid. If barium nitrate is used, it should be followed by dilute nitric acid.
The acids are added to distinguish between a sulphate and a carbonate. Both barium sulphate and barium carbonates are insoluble and give what precipitates. A carbonate would dissolve in dilute hydrochloric acid or nitric acid. A sulphate odes not dissolve in either of these acids.
The table below summarizes the process for any sulphate.
Test Observation Reaction taking place
To a solution of sulphuric acid in a test tube, add dilute hydrochloric acid or nitric barium chloride or barium nitrate solution. White precipitate of barium sulphate is formed which is insoluble in the acid BaCl_(2 (aq))+ H_2 SO_(4 (aq)) →BaSO_(4 (S))+ 2HCl_( (aq))
Or
Ba(NO_3 )_2+ H_2 SO_4(aq) →BaSO_(4 (s) )+ 2HNO_(3 (aq))
To copper sulphate solution in a test tube, add dilute hydrochloric acid White precipitate formed, insoluble in the acid → CuSO_(4 (aq))+ BaCl_(2 (aq))→BaSO_(4 (s))+ CuCl_(2 (aq))
By use of barium chloride or barium nitrate, any soluble sulphate can be detected.
Hydrogen sulphide
Preparation of hydrogen sulphide
By action of dilute hydrochloric acid on iron(II) sulphide
Caution!!
This gas is poisonous and should be prepared in a fume cupboard.
Equation of reaction
Iron(II) sulphide + hydrochloric acid → iron(II) chloride + hydrogen sulphae
FeS_((s)) + 2HCl_( (aq)) → FeCl_(2(s)) + H_2 S_( (g))
The gas is collected over warm water since it is soluble in cold water. If required dry. It can be passed over undhydrous calcium chloride.
Dilute hydrochloric acid may be used to dry the gas instead of calcium chloride.
Preparation of hydrogen sulphide
Properties of hydrogen sulphide
Physical properties
Colourless poisonous gas
It is slightly denser than air.
Has a repulsive smell – characteristic of rotten eggs. (proteins in eggs produce hydrogen sulphide when they decay).
Has a boiling point of 620C (211K).
It is fairly soluble in water for form a weak acidic solution and only slight ionization takes place.
H_2 S_((g)) + H_2 O_((l)) ⇌ S_((aq))^(2-) + 2H_2i O_((aq))^+
Chemical properties of hydrogen sulphide
Burning in oxygen
When a solution of hydrogen sulphide is left exposed to air for a few days, a white precipitate of amorphus, sulphur is formed.
2H_2 S_( (aq)) + O_(2 (g)) → 2H_2 O_( (l)) + S_( (s))
It burns with a pale blue flame to form sulphur(IV) oxide and steam.
2H_2 S_( (aq)) + 3O_(2 (g)) → 2H_2 O_((g)) + 2SO_(2 (g))
Indirect sunlight, the two reactions above (I) and (II) take place simultaneously and the equations are as follows:
2H_2 S_((g)) + 2O_(2 (g)) → 2H_2 O_((l)) + SO_(2 (g)) + S_((s))
Reducing actions
Hydrogen sulphide is a powerful reducing agent in all these reactions; it supplies electrons producing a yellow precipitate of sulphur.
Reaction with concentrated nitric acid
It reduces nitric acid to nitrogen(IV) oxide and itself oxidizes to sulphur.
2HNO_(3 (aq)) + H_2 S_( (g)) → S_( (g)) + 2H_2 O_( (l)) + 2NO_(2 (g))
Reaction with concentrated sulphuric acid
It reduces sulphuric acid to water and sulphur is deposited.
3H_2 S_( (g)) + H_2 SO_(4 (l)) → 4S_( (s)) + 4_2 O_( (l))
Reduction of iron(III) chloride
It reduces iron(III) chloride to iron(II) chloride
2FeCl_(3 (aq)) + H_2 S_((g)) → 2FeCl_(2 (aq)) + 2HCl_( (aq)) + S_( (s))
Yellow or brown green
Ionically, the above reaction can be written as;
2Fe_(3 (aq))^+ + S_((aq))^(2-) → 2Fe_((aq))^(2+) + S_((s))
Reduction of potassium manganese(VII) and potassium dichromate(VII)
The potassium manganese(VII) solution becomes decolourized as the manganese(VII) ions are reduced to manganese(II) ions.
2KMnO_(4 (aq)) + 3H_3 SO_(4 (aq)) + 5H_2 S → K_2 SO_(4(aq)) + 2MnSO_(4 (aq)) + 8H_2 O_( (l))
Ionically
2MnO_( 4 (aq))^- + 5H_2 S_( (g)) + 〖6H〗_((aq))^+ → 2Mn^(2+) + 8H_2 O_((l))
Purple colourless
The dichromate(VI) ions are reduced to chromate(III) ions. The solution changes from orange to green.
K_2 Cr_2 O_(7(aq)) + 3H_2 S_((g)) + 4H_2 SO_4 → K_2 SO_(4(aq)) + 5S_((s)) + 4H_2 O + Cr_2 (SO_4 )_(3(aq))
Orange green
Reduction of bromine water
The bromine water changes from reddish-brown to colourless. This bromine water is reduced to bromide ions, sulphur is precipitated.
Br_(2(aq)) + H_2 S_((g)) → 2HBr_((aq)) + S_((s))
Reddish brown colourless
Reduction of hydrogen peroxide solution
Hydrogen peroxide is reduced to water and sulphur is deposited.
H_2 O_(2(aq)) + H_2 S_((g)) → S_2 + 2H_2 O_((l))
Action on metallic salts
Metallic sulphides are precipitated
With lead nitrate, a black precipitate of lead sulphide is precipitated.
Pb(NO_3 )_2 + H_2 S_((g) ) → PbS_((s) ) + 2HNO_(3 (aq))
Black precipitate
With copper(II) sulphate
A dark brown precipitate of copper(II) sulphide is formed.
CuSO_(4(aq)) + H_2 S_((g)) → CuS_((s)) + H_2 SO_(4(aq))
Action of hydroxides
Hydrogen sulphate acts a weak dibasic acid
H_2 S_((g)) ⇌ 2H_((aq))^+ + S_((aq))^(2-)
2NaOH_((aq)) + H_2 S_((g)) → Na_2 S + H_2 O_((l))
Sodium sulphide
In excess hydrogen sulphide; sodium hydrogen sulphide is formed.
NaOH_((aq)) H_2 S_((g)) → NaHS_((aq)) + H_2 O_((l))
Pollution effects of sulphur compounds to the environment
Pollution is a change in the conditions of the environment such that the life of the organisms in the environment become unsafe. This change is often brought about by the addition of the substance to the environment at a faster rate than the environment can cope with.
A substance that pollutes the environment is called a pollutant. It must have a harmful effect before it can be considered to be a pollutant.
The compounds of sulphur that are pollutants; sulphur dioxide, sulphur(VI) oxide, sulphurous and sulphuric acid and hydrogen sulphide
The sulphur compounds affect the lungs, villi, lichens and other plants and attack fabrics, masory (buildings and metals)
These compounds of sulphur are released to the environment from industries and manufacturing process by burning sulphur or sulphide ores, for example in the extraction of metals like copper, lead, zinc and manufacture of sulphuric acid. Burning of sulphur or sulphide ores such as zinc blende (ZnS), galena (PbS), copper pyrites (CuFeS_2), iron pyrites (FeS_2), all emit sulphur compounds in the atmosphere.
QUESTIONS
(a) (i) Name one substance that is reacted with hydrochloric acid to produce sulphur dioxide in the laboratory.
(ii) State the conditions for the reactions
(iii) Name the substance that can be used to dry the sulphur dioxide formed.
(iv) Write equation for the reaction leading to the formation of sulphur dioxide.
(b) State what would be observed and explain what would happen if sulphur dioxide is passed through a solution containing
(i) acidified potassium dichromate
(ii) dye
(c) Briefly describe how sulphur dioxide can be converted to sulphuric acid. Your answer should include equations and conditions for the reaction.
(Qn. 11, UNEB, 2002)
(a) Write an equation for the formatin of sulphur dioxide was bubbled through an acidified solution of potassium dichromate
(b) Sulphur dioxide was bubbled through an acidified solution of potassium dichromate.
(i) State what was observed.
(ii) Briefly explain your answer in (b) (i). (Qn. 9, UNEB, 2005)
(a) (i) State the conditions under which sulphuric acid reacts with ethanol to form ethane
(ii) Write an equation for the formation of ethane from ethanol and sulphuric acid.
(iii) State the property of sulphuric acid shown in the reaction in (a) (ii)
(b) Concentrated sulphuric acid reacts with graphite according to the following equation.
C_2 + H_2 SO_(4 (l)) → CO_(2(g)) + 2SO_(2(g)) + 2H_2 O_((l))
Calculate the mass of carbon that can react completely with a solution containing 19.6g of sulphuric acid. (UNEB)
(a) Sulphur dioxide can be prepared by roasting zinc blende in air according to the following equation
2ZnS_((s)) + 3O_(2(g)) → 2SO_(2(g)) + 2ZnO_((s))
Calculate the volume of sulphur dioxide evolved at room temperature when 9.7g of zinc sulphide is reacted with excess oxygen Zn=65,S=32; 1 mole of gas occupies 24dm^3at room temperature.)
(b) During the manufacture of sulphuric acid by contact process, sulphur dioxide is heated with oxygen in the presence of a catalyst.
(i) Name the crystal
(ii) Write the equation of reaction between sulphuric acid and oxygen
(Qn. 10, UNEB, 2017)
(a) (i) State the conditions under which sulphuric acid reacts with sodium nitrate to form nitric acid.
(ii) Write equation for the reaction in (a) (i) above
(b) Sulphur was warmed with concentrated nitric acid
(i) State what was observed
(ii) Write equation for the reaction (Qn. 6, UNEB, 2007)
(a) (i) With the aid of a labeled diagram, explain how a pure dry sample of sulphur dioxide can be prepared in the laboratory. Using sodium sulphite and sulphuric acid.
(ii) Write an equation leading to the formation of sulphur dioxide
(b) Name one reagent that would confirm the presence of sulphur dioxide and state what would be observed if the reagent you have named was threatened with sulphur dioxide.
(c) Write an equation to show the reaction between sulphur dioxide and
(i) water
(ii) Oxygen in the presence of hot platinum
(d) The product of reaction in (c) (ii) was mixed with water and barium nitrate solution added to the resultant mixture.
(i) State what was observed.
(ii) Explain what took place (equation not required) (Qn. 11, UNEB, 2014)
CHLORINE AND ITS COMPOUNDS
General objective
By the end of this topic, the learners should be able to explain the Chemistry of chlorine and its compounds.
Specific objectives
Learners should be able to:
Describe and explain the laboratory preparation and manufacture of chlorine.
Outline the properties of chlorine.
Outline the uses of chlorine.
Sub-topic
Hydrogen chloride
Learners should be able to:
Explain preparation of hydrogen chloride in the laboratory.
Explain how the composition of hydrogen chloride can be reduced from a series of chemical reactions.
Explain the properties of hydrogen chloride.
Explain the effect of a solvent on the properties of hydrogen chloride.
State the uses of hydrogen chloride.
Test for chlorine ions in the laboratory.
State the uses of hydrochloric acid.
Chlorine
Is a number of group (VII) element which are called halogens. Other members of group (VII) are Fluorine(F), Bromine (Br) and Iodine (I)
Chlorine has atomic number 17 and electronic arrangement (configuration) 2:8:7
It attains the noble gas electronic structure by gaining or sharing electrons.
E.g. By electrons, it forms negatively charged ion, the chloride ions and attains the structure of the gas Argon.
Cl + e → Cl^-)
2:9:7 2:8:8
By gaining electrons, it can form ionic compounds e.g. sodium chloride.
By sharing electrons, it can form covalent compound (or covalent bonds)
E.g. Cl-Cl or H-Cl
Or
+→
Methods of preparation of chlorine
Laboratory preparation
(i) From potassium manganate(VII) and concentrated hydrochloric acid.
Cation!! Chlorine is poisonous and should be prepared in a fume cupboard.
There is no hearing during the reaction.
Laboratory preparation chlorine
The concentrated hydrochloric acid is oxidized by potassium manganate(VII) crystals. Effervescence takes place in the cold as the gas is evolved.
Equation of reaction
2KMnO_(4 (s)) + 16HCl_((aq)) → 2KCl_((aq)) + 2MnCl_(2 (aq)) + 8H_2 O_((l)) + 5Cl_(2 (g))
The gas is passed through water to remove unreacted hydrogen chloride. It is dried by concentrated sulphuric acid. It is collected by downward delivery.
2. Preparation of chlorine from manganese(IV) oxide and concentrated hydrochloric acid
The mixture is heated. (This experiment must be carried out in a fume cupboard as the gas is poisonous. Effervescence take place. A greenish – yellow gas is evolved, which together with a certain amount of hydrogen chloride passes into the first bottle containing water. This removes hydrogen chloride and the concentrated sulphuric acid in the second bottle dries the gas.
Equation of reaction
MnO_(2 (s)) + 4HCl_((aq)) → MnCl_(2(aq)) + 2H_2 O_((l)) + Cl_(2 (g))
Experimental set up for preparation of chlorine
3. Preparation of chlorine from bleaching powder
Bleaching powder is reacted with hydrochloric acid or nitric acid. Effervescence takes place and a greenish – yellow gas is collected by the method shown for the preparation from potassium manganate(VII). Heat is not required.
Equation of reaction
CaOCl_(2(s)) + 2HCl_((aq)) → CaCl_(2(aq)) + H_2 O_((l)) + Cl_(2(g))
Likewise
CaOCl_(2(s)) + 2HNO_(3(aq)) → Ca(NO_3 )_2(aq) + H_2 O_((l) ) + Cl_(2(g))
4. By heating concentrated sulphuric acid with a mixture of any chloride and manganese(IV) oxide
E.g. 2NaCl_((s)) + MnO_(2(s)) + H_2 SO_(4(l)) → Na_2 SO_4 + H_2 O_((l)) + Cl_((g))
Note: In this reaction, the sodium chloride first reacts with concentrated sulphuric acid to produce hydrogen chloride which is then oxidized by manganese(IV) oxide to produce chlorine.
Industrial preparation of chlorine
Chlorine is prepared on a large scale (commercially) by electrolysis of strong sodium chloride solution (brine). In the process both chlorine and sodium hydroxide are produced. Since chlorine and sodium hydroxide react with each other, they are kept separate from each other. In the mercury cathode cell, a set of graphite rods dipping into the brine serves as the anode.
The base of the tank is filled with mercury which serves as the cathode. The mercury circulates slowly through the electrolytic cell.
At the Cathode
Sodium is discharged preferentially to hydrogen and dissolves in mercury to form a sodium amalgam (Na/Hg). Sodium amalgam is passed through a trough below the cell which contains water. Here it reacts to form a solution of sodium hydroxide and hydrogen is liberated. The mercury is regenerated and used again.
2Na/Hg_((g) ) + 2H_2 O_((l) ) → 2NaOH_((aq))^+ + Hg_((l) ) + H_(2(g))
At the anode
The chloride ions from the brine (sodium chloride solution) are discharged as chlorine gas.
2Cl_((aq))^- – 2e → 2Cl_(2(g))
The mercury cathode cell
Summary of electrode reactions
At cathode (negative electrode)
Na_((aq))^+ + e → Na_((s))
At the anode (positive electrode)
Cl_((aq))^- → Cl_((g)) + e
Then: two chlorine atoms combine
2Cl_((g)) → Cl_(2(g)) Or
2Cl^- → Cl_(2(g)) + 2e
Physical properties of chlorine
Greenish-yellow gas with unpleasant irritating, chocking smell.
It is very poisonous if inhaled to even a small extend.
It is about 2½ times denser than air
Slightly soluble in water forming chlorine water (hydrochlorous acid)
Chemical properties of chlorine
Does not burn or support combustion.
React with most metals or non-metals to form corresponding chlorides.
Reaction with metals
Burning magnesium and sodium continue to burn in chlorine forming a white solid.
E.g. 2Na_((s)) + Cl_(2(g)) → 2NaCl_((s))
Mg_((s)) + Cl_(2(g)) → MgCl_(2(s))
Reaction with “Dutch metal”, an alloy of copper and zinc
A thin sheet of Dutch metal burns in chlorine with a green flame forming copper and zinc chlorides.
i.e.Cu_((s)) + Cl_(2(g)) → CuCl_(2(s))
Zn_((s)) + Cl_(2(g)) → ZnCl_(2(s))
The green flame is due to copper
Iron reactions with chlorine to form iron(III) chloride
2Fe_((s)) + 3Cl_(2(g)) → 2FeCl_(3(s))
Iron(II) is not formed because it is immediately oxidized by chlorine to iron(III) chloride.
Note:
Reaction of iron with HCl gives iron(II) chloride plus hydrogen gas.
Aluminium also reacts with chlorine to form aluminium chloride
2Al_((s)) + 3Cl_(2(g)) → 2AlCl_(3(s))
Reaction with non-metals
Phosphorus – burns spontaneously in chlorine forming white fumes of phosphorus.
Chloride and phosphorus(V) chloride
i.e. 2P_((s)) + 3Cl_(2(g)) → 2PCl_(3(g))
2P_((s)) + 5Cl_(2(g)) → 2PCl_(5(g))
Action of chlorine on sulphur
When chlorine is passed over heated molten sulphur in a distillation flask, connected to a condenser, a reddish liquid of sulphur(I) chloride distils over.
2S_((l)) + Cl_(2(g)) → S_2 Cl_(2(l))
H_2 S_((g)) + Cl_(2(g)) → S_((s)) + 2HCl_((g))
With sulphites
Sulphites are oxidized by chlorine to sulphates.
Cl_(2(g)) + H_2 O_((l)) + Na_2 SO_(3(aq)) → 2HCl_((aq)) + Na_2 SO_(4(aq))
Action of chlorine on sulphorous acid
Sulphorous acid is oxidized to sulphuric acid. Sulphorous is formed when sulphur(IV) is dissolved in water.
2H_2 O_((l)) + SO_(2(g)) + Cl_(2(g)) → H_2 SO_(4(aq)) + 2HCl_((aq))
The presence of the sulphuric acid can be shown by testing with dilute hydrochloric acid and barium chloride. A white precipitate of barium sulphate is obtained.
Reaction of chlorine with alkalis
Calcium hydroxide
When chlorine is passed over calcium hydroxide bleaching powder (calcium hypo chloride) is formed. The colour of chlorine disappears immediately.
Ca(OH)_2(s) + Cl_2(g) → CaOCl_2∙2H_2 O_((s))
Calcium hydroxide calcium bleaching powder
Sodium and potassium hydroxides
These have similar reactions with chlorine and react in two ways i.e. pending on the concentration and temperature of the alkali used.
In cold and dilute solutions
2NaOH_((aq)) + Cl_(2(g)) → NaCl_((aq)) + NaClO_((aq)) + H_2 O_((l))
Sodium chloride Sodium chloride(I)
When hot concentrated solution of the alkali is used, iE is possible the chlorate(I) is first formed
i.e. Cl_(2(g)) + NaOH_((aq)) → NaOH_((aq)) + NaOHCl_((aq)) + H_2 O
Then the chlorate(I) is converted to chlorate(V) and a chloride.
3NaClO_((aq)) → 2NaCl_((aq)) + NaClO_(3(aq))
Sodium chloride(V)
Reaction with iron(II) chloride
When chlorine is passed through iron(II) chlorine, the colour of the solution turns from green to reddish-brown.
Reason
Chlorine oxidizes iron(II) chloride (green) to iron(III) chloride (reddish-brown)
2FeCl_2 + Cl_2 → 2FeCl_3
green reddish brown
Reaction with iron
When chlorine is passed over heated iron wire in the apparatus below, the iron glows red hot forming iron(III) chloride which sublimes and condenses on the colder parts to give a black solid.
2Fe_((s)) + 3Cl_(2(g)) → FeCl_3
This reaction can serve as a method for preparing iron(III) chloride.
Note:
The red hot glow of iron shows that the reaction produces a lot of heat (exothermic)
Iron(II) chloride is not formed but iron(III) chloride because chlorine oxides iron to its lightest (oxidation state) charge Fe^(3+).
If hydrogen chloride is used instead of chlorine iron(II) chloride would be formed instead and the excess gas would consist of unreacted hydrogen chloride and hydrogen gas formed by the reaction.
2HCl_((g)) + Fe_((s)) → FeCl_2 + H_(2(g))
Preparation of iron(III) chloride
N.B: The bleach crystals of anhydrous iron(III) chloride should be removed and placed in a discator, as they are deliquescent (absorb) water vapour from the atmosphere and dissolve in it to form a solution (brown solution).
Reaction with sodium sulphite
Chlorine oxidize a solution of sodium sulphite to sodium sulphate by the reaction.
Cl_(2(g)) ¬ + H_2 O_((l)) + Na_2 SO_(4(aq)) + HCl_((aq))
Note:
The presence of the sulphate can be confirmed by adding barium chloride or barium nitrate to the solution, giving a white precipitate.
i.e. SO_(4(aq))^(2-) + Ba_((aq))^(2+) → BaSO_(4(s))
white precipitate
Reaction with hydrogen sulphide
When a gas jar of chlorine is inverted over one of hydrogen sulphide, a yellow deposit of sulphur is left.
H_2 S_((g)) + Cl_(2(g)) → 2HCl_((g)) + S_((s))
Reaction with hydrogen
Hydrogen burns in chlorine to give hydrogen chloride. Chloride has a high affinity for hydrogen.
H_(2(g)) + Cl_(2(g)) → 2HCl_((g))
Reaction with turpentine (hydrocarbon)
When a piece of turpentine is dipped into a warm gas jar of chlorine, a violent reaction occurs with red flash and black deposit of carbon are observed.
The turpentine combines violently with chlorine forming hydrogen chloride gas leaving behind black deposit of carbon.
i.e. C_10 H_16 + 8Cl_(2(g)) → 10C_((s)) + 16HCl_((g))
turpentine
Displacement reactions of chlorine
The reactivity of the halogens decrease down the group in the order.
F_2>Cl_2>Br_2>I_2
Reaction with potassium bromide
2KBr_((aq)) + Cl_(2(g)) → 2KCl_((aq)) + Br_(2(g))
When chlorine is bubbled through a solution of potassium bromide, a reddish-brown solution is formed.
Reason
Chlorine is higher than bromine in the electrochemical series and displaces the bromide ion from its solution forming reddish-brown bromine.
2Br_((aq))^- + Cl_(2(g)) → Br_2(aq or l) + Cl_((aq))^-
Reaction with potassium iodide solution
Passing chlorine through a solution of potassium iodide solution gives a brown solution.
2Kl_((aq)) + Cl_(2(g)) → 2KCl_((aq)) + I_(2(aq))
deep brown solution
Or
2I_((aq))^- + Cl_(2(g)) → I_(2(aq)) + 2Cl_((aq))^-
Explanation for the observation is as above.
Effect of sunlight on chlorine water
Uses of chlorine
As a disinfectant in the treatment of water works, sewage and swimming pools.
As a bleaching agent. It is used to bleach cotton, silk, onen and wood pulp. It is not used to bleach silk and wood because it destroys them.
Also used in the manufacture of bleaching powder like calcium hypochlorite by passing chlorine gas in calcium hypochlorite by passing chlorine gas in calcium hydroxide.
Ca(OH)_2(aq) + Cl_2(g) → CaOCl_2 + H_2 O_((l))
calcium hypochlorite
In the manufacture of organic compounds.
e.g. Carbon tetrachloride or tetra chloromethane CCl_4, used as a fire extinguisher and also as a solvent.
Chloromethane CH_3 Cl (Chloroform) used as an anesthetic
Manufacture of plastics e.g. PVC (polyvinyl chloride), insecticides and many other organic compounds.
Manufacture of weed killers e.g. sodium chlorate
Manufacture of DT (Dinitrotri Chloromethane)
Manufacture of antiseptics e.g. detol.
Manufacture of trichloroethane C_2 H_3 Cl_3 a solvent for removing greases from clothes.
Hydrogen chloride
Preparation of hydrogen chloride from sodium chloride
It is prepared by the action of concentrated sulphuric acid on common salt (sodium chloride or any chloride).
The gas is collected by passing it through a wash bottle of concentrated sulphuric acid to dry it.
It is collected by downward delivery since it is denser than air.
The reaction proceeds in the cold; heating is required.
Laboratory preparation of hydrogen chloride
Equation of reaction
NaCl_((s)) + H_2 SO_(4(l)) → NaHSO_(4(s)) + HCl_((g))
Or
H_2 SO_(4(aq)) + Cl_((aq))^- → HSO_((aq))^- + HCl_((g))
In this reaction, a volatile acid (HCl) is being displaced from its salt by a less volatile acid (H_2 SO_4).
NaHSO_(4(s)) + NaCl_((s)) → Na_2 SO_(4(aq)) + HCl_((g))
This reaction shows the acidic nature of sodium hydrogen sulphate.
Physical properties of hydrogen chloride
Colourless gas with irritating, chocking smell.
It emits white fumes in air because it dissolves tiny drops of hydrochloric acid.
HCl_((g)) + H_2 O_((l)) → H_3 O_((aq))^+ + Cl_((aq))^-
It is very soluble in water. 1cm3 of water dissolves in 450cm3 of hydrogen chloride. This can be shown by a fountain experiment.
It turns moist blue litmus paper red (acidic)
It reacts with ammonia to form dense white fumes of ammonia chloride (test for HCl).
NH_(3(g)) + HCl_((g)) → NH_4 Cl_((s))
Ammonium chloride forms dense white fumes
Does not burn and extinguishes a burning splint.
It is 1¼ times denser than air.
Its boiling and melting points are 188K(-850C) and 159K(-1140C) respectively.
Chemical properties
Reaction with ammonia
Forms dense white fumes of ammonium chloride
NH_(3(g)) + HCl_((g)) → NH_4 Cl
Ammonium chloride seen as white fumes
Reaction with iron
Passing hydrogen chloride over heated iron wire gives green iron(II) chloride.
Fe_((s)) + 2HCl_((g)) → FeCl_(2(s)) + H_(2(g))
Iron(II) chloride (green solid)
Forms a white precipitate with acidified silver nitrate solution
AgNO_(3(aq)) + HCl_((aq)) → AgCl_((S)) + HNO_(3(aq))
Ag_((aq))^+ + Cl_((aq))^- → AgCl_((s))
The precipitate dissolves in aqueous ammonia forming a colourless solution. This reaction is used to confirm the chloride ion.
AgCl_((S)) + 2NH_(3(aq)) → Ag(NH_3 )_2^+ + Cl_((aq))^-
Hydrochloric acid
Hydrogen chloride dissolves in water to form hydrochloric acid. But due to its high solubility, it should not be dissolved in water directly. Use a funnel to avoid “suck back.”
The experimental set up can be like that of preparing hydrogen chloride gas but replace the gas jar with a funnel arrangement as shown below. This solution is now hydrochloric acid.
Physical properties of hydrochloric acid
Hydrochloric acid whether dilute or concentrated behave almost the same unlike sulphuric acid and nitric acid.
Dilute hydrochloric acid has a sour taste and turns blue moist litmus paper red.
Concentrated hydrochloic acid is a colourless fuming liquid due to the escape of hydrogen chloride gas in the air. It dissolves in the water to form tiny drops of hydrochloric acid.
Chemical properties
Reacts with metals above hydrogen in the electrochemical series to give hydrogen gas.
Zn_((s)) + 2HCl_((aq)) → ZnCl_(2(aq)) + H_(2(g))
Ionically
Zn_((s)) + 2H_((aq))^+ → Zn^(2+) + H_(2(g))
Reacts with carbonates and hydrogen carbonaes to give carbondioxide gas.
Na_2 CO_(2(s)) + 2HCl_((aq)) → 2NaCl_((aq)) + CO_(2(g)) + H_2 O_((l))
Ionicially
CO_(3(aq))^(2-) + 2H^+ → CO_(2(g)) + H_2 O_((l))
It reacts with alkalis and basic oxides to give salt and water only.
NaOH_((aq)) + HCl_((aq)) → NaCl_((aq)) + H_2 O_((l))
CaO_((s)) + 2HCl_((aq)) → CaCl_2 + H_2 O_((l))
CaO_((s)) + 2H_((aq))^+ → Ca_((aq))^(2+) + H_2 O_((l))
It is oxidized to chlorine by potassium manganate(VII) and manganese(IV) oxide
2KMnO_(4(s)) + 16HCl_((aq)) → 5Cl_(2(g)) + 2MnCl_(2(aq))
The above reaction does not need any external source of heat.
MnO_(2(g)) + 4HCl_((aq)) → Cl_(2(g)) + MnCl_(2(aq)) + 2H_2 O_((l))
This reaction needs heating (refer to preparation of chlorine).
Hydrochloric acid dissolved in methyl benzene (toluene) has different properties from that dissolved in water (aqueous one).
In water, hydrochloric acid, being a strong acid is completely ionized as: –
HCl_((aq)) → H_((aq))^+ + Cl_((aq))^-
Or
HCl_((aq)) + H_2 O_((l) ) → H_3 O_((aq))^+ + Cl_((aq))^-
This is because water is a polar solvent and the ions are hydrated (surrounded by water molecules). This exhibits its acid properties.
In methyl benzene (toluene) a non-polar solvent, hydrogen chloride exists in a molecular state. It forms no ions. As such, it exhibits the following properties in toluene:
It does not turn moist blue litmus red. This is because no hydroxonium ions (H_3 )^+ ) are present.
It does not react with carbonates and hydrogen carbonates.
Does not conduct electricity.
Does not produce hydrogen with zinc and magnesium.
It forms a white precipitate with ammonia. This is because ammonia chlorine is insoluble in toluene.
All the above observations indicate that hydrogen chloride in methyl benzene (toluene) is in molecular form and not ionized.
Chlorides
Preparation of chlorides
Soluble chlorides
Those of potassium, sodium and copper are prepared by reacting the oxide of the metal or carbonate with dilute hydrochloric acid.
Na_2 O_((s)) + 2HCl_((aq)) → 2NaCl_((aq)) + H_2 O_((l))
CaCO_(3(s)) + 2HCl_((aq)) → CaCl_(2(s)) + CO_(2(g)) + H_2 O_((l))
CuO_((s)) + 2HCl_((aq)) → CuCl_2 + H_2 O_((l))
Chlorides of magnesium, aluminium, zinc and iron are prepared by either reacting the metal or its oxide or carbonates with dilute hydrochloric acid.
Zn_((s)) + 2HCl_((aq)) → ZnCl_(2(aq)) + H_2 O_((l))
Al_2 O_(3(s)) + 6HCl_((aq)) → 2AlCl_(3(aq)) + 3H_2 O_((l))
Note:
Aluminium chloride in vapour from exists as a dimer; Al_2 CL_6.
Its sodium in water is acidic because of the reaction.
AlCl_(3(s)) + 3H_2 O_((l)) → Al(OH)_3(s) + 3HCl_((aq))
The HCl_((aq)) is a strong acid and ionizes complete as: HCl_((aq)) → H_((aq))^+ + Cl_((aq)). This makes the solution acidic.
Aluminium chloride sublimes on heating.
Iron(III) chloride is made by passing chlorine over heated iron wire. Its solution in water is also acidic.
FeCl_(3(aq)) + 3H_2 O_((l)) → Fe(OH)_3(s) + 3HCl_((aq))
Insoluble chlorides
Those of silver and lead are prepared by precipitation reaction (double decomposition).
A soluble salt of lead or silver is added to a solution containing a chloride ion. i.e. Pb_((aq))^(2+) + 2Cl_((aq))^- → PbCl_(2(s))
e.g. Pb(NO_3 )_2(aq) + 2NaCl_((aq) ) → PbCl_2(s) + 2NaNO_(3(aq))
Ag_((aq))^+ + Cl_((aq))^- → AgCl_((s0)
Preparation of unhydrous chloride
By passing chlorine or hydrogen chloride over heated metal.
Unhydrous iron(III) chloride
2Fe_((s)) + 3Cl_(2(g)) → 2FeCl_(3(s))
Iron(III) chloride is collected as a black sublimate.
Unhydrous iron(II) chloride is prepared by passing hydrogen chloride over heated iron wire.
Fe_((s)) + 2HCl_((aq)) → FesCl_(2(s)) + H_(2(g))
Test for chlorides
Chlorides in the presence of nitric acid give a white precipitate of silver nitrate.
Ag_((aq))^+ + Cl_((aq))^- → AgCl_((g))
Lead nitrate solution – is sometimes used instead of silver nitrate solution. A white precipitate which dissolves on heating and reappears on cooling confirms the presence of a chloride.
Pb_((aq))^(2+) + 2Cl_((aq))^- → PbCl_(2(s))
Solubility of chlorides
To any solid chloride, add concentrated sulphuric acid. Hydrogen chloride gas is evolved. The gas forms white fumes with ammonia and a white precipitate with acidified silver nitrate solution.
Reactions taking place
NaCl_((s)) + H_2 SO_(4(l)) → NaHSO_(4(s)) + HCl_((g))
HCl_((g)) + NH_(3(g)) → NH_4 Cl_((s))
White fumes
Or
HCl_((g)) + AnNO_(3(aq)) → AgCl_((s)) + HNO_3
White precipitate
Questions
(a) Chlorine can be prepared in the laboratory using potassium manganate(VII) to produce chlorine.
Name one substance that can react with potassium manganate(VII) to produce chlorine.
State the conditions for the reaction
Write an equation leading to the formation of chlorine.
(b) Damp blue litmus paper was dropped into a gas jar of chlorine. State what was observed. Explain you observations.
(c) A boiling tube filled with chlorine was inverted into a beaker containing chlorine water and exposed to sunlight for some time.
(i) State what was observed
(ii) Explain with the aid of equation(s), your observations in (c) (i)
(d) Write an equation to show how chlorine can react with
(i) dilute potassium hydroxide solution
(ii) turpentine C_10 H_16
(e) Briefly describe a test you would carry out to confirm the presence of a chloride ion in solution. State what would be observed and write equation for the reaction that would take place. (Qn. 13, UNEB, 2014)
(a) Describe how a pure dry sample of chlorine can be prepared in the laboratory from potassium (manganate(VII) crystals. (Your answer should include a well labeled diagram and equation for the reaction.)
(b) State what would be observed if chlorine is passed through a
(i) blue litmus solution
(ii) Potassium bromine solution
(iii) Solution of iron(II) ions
(c) Write equation for the reaction in (b) (ii) and (iii)
(d) Write equatin for the reaction between chlorine and
(i) heated ion
(ii) cold dilute sodium hydroxide (Qn. 11, UNEB, 2014)
(a) Describe how a pure sample of chlorine can be prepared in the laboratory starting from potassium manganate(VII). (No diagram is required, but your description should include conditions and equations for the reaction.)
(b) State what would be observed and write equation(s) for the reaction(s) that would occur if
(i) chlorine was bubbled into an aqueous sodium hydroxide
(ii) burning magnesium was lowered into a gas jar of dry chlorine.
(iii) chlorine ions passed through a solution of potassium iodide solution.
(Qn. 14, UNEB, 2012)
(a) State the conditions under which iron can react with hydrochloric acid and write equation for the reaction.
(b) Draw a fully labeled diagram for the set up of the apparatus which
(c) A student left a slasher made of iron on the compound for two weeks. State what was observed and explain your answer.
(d) (i) Name one reagent that can be used to distinguish between iron(II) sulphate and iron(III) sulphate.
(ii) State what would be observed if the reagent you have named in (d) (i) above were separately treated with two iron salts and write equation for the reactions. (Qn. 12, UNEB, 2012)
(a) Hydrogen chloride can be prepared from sodium chloride according to the following ionic equation.
Cl_((aq))^- + H_((aq))^+ → HCl_((g))
Calculate the mass of sodium chloride that would be required to produce 3.60dm3 of hydrogen chloride measured at room temperature.
(H=1,Na=23,Cl=35.5,one mole of a gas at room temperature occupies 24.0dm^3
(b) State what would be observed and in each case, write an equation for the reaction that would take place when
(i) an aqueous solution of hydrogen chloride was added to solution containing lead(II) ions(II)
(ii) excess hydrogen chloride was passed over strongly heated iron wire.
(c) Briefly explain the following observations and in each case, illustrate your answer with equation(s).
(i) anhydrous iron(II) chloride cannot be prepared by direct synthesis using iron and chlorine.
(ii) an aqueous solution of hydrogen chloride gives a white precipitate with silver nitrate whereas a solution of hydrogen chloride in tetra chloromethane shows no observable change when treated with silver nitrate solution. (Qn. 11, UNEB, 2009)
APPLIED CHEMISTRY
General objective
By the end of this topic, the learners should be able to appreciate the importance and application of Chemistry in everyday life.
Sub topic
Mineral resources and industrial processes.
Specific objective
Learners should be able to:
Outline the application of electrolysis of sodium extraction.
Describe how iron is extracted by reduction.
Outline how copper is refined by electrolysis.
Define an alloy.
State some common alloys and give their composition.
Describe how sugar is manufactured in an industry.
Extraction of metals
Apart from silver, gold and platinum which are found in free or pure state, most metals are found as minerals (compounds of the metal) mixed with earthly material (gangiies). Such a mixture is called an ore.
The source of the metals and the method used to extract them depends on their nature.
General properties of metals
Metals are elements that ionize by loss of electrons to form positively charged ions.
e.g. K_((g)) → K_((aq))^+ + e
Ca_((g)) → Ca_((g)) + 2e
Al_((g)) → Al^(3+) + 3e
The number of electrons lost per atom of the metal is equal to the valency of the metal. Most metals have 1 – 3 electrons more than a noble gas.
Metals therefore react by loss of electrons forming ionic compounds.
Physical of metals
Mallable – can be beaten into sheets
Ductile – can be drawn into wire
Lustrous – can be polished and shine
Strong and tough – have high tensile strength and can stretch.
Good conductors of electricity and heat.
High density (e.g. mercury 13.6g/cm3)
Exceptional physical properties
Mercury is a liquid (m.p – 390C)
The metals sodium and potassium have low densities 0.97 and 0.86gcm-3 and float on water. Their melting points are lone (Na 980C, K 630C) and they are soft enough to be cut by pen knife.
Solid metals consist of closely parked ions (cations) held together by mobile electrons. It is difficult to separate the ionic structure and their melting and boiling points are high.
Many metals consist of layers of ions which can easily slip over each other. Therefore, metals are malleable and ductile. The mobile electrons in metals can carry an electric current and also heat energy metals are therefore good conductors.
Chemical properties of metals
Burn in oxygen to form basic oxides.
React with acids to form salts.
e.g. Al_((s)) + 6HCl_((aq)) → 2AlCl_3 + 3H_(2(g))
3Pb_((s)) + 8HNO_(3(aq)) → 3Pb(NO_3 )_2(aq) + 4H_2 O_((l) ) + 2NO_((g))
Do not usually react with hydrogen.
Their chlorides are ionic (i.e. non-volatile electrolytes soluble in and not hydrolyzed by water.
Reducing agents.
Metals at their electrons easily. This they are reducing agents.
e.g. Zn_((s)) → Zn^(2+) + 2e
Zn_((s)) + Cl_(2(g)) → ZnCl_2
Zinc metal reduces chlorine to chloride ions and is itself oxidized.
Preliminary treatment of ores
Most metals are obtained as ores containing varying amounts of impurities. Some of these impurities can be separated out before the final extraction of the metal. Metal ores can be treated by means of the following methods or a combination of methods.
Washing – materials live soil can be removed by washing the ore with plenty of clean water.
Concentration – (froth – floatation) – the crashed ore is mixed with water containing a detergent or suitable solvent.
Examples sodium hydroxide is used for bauzate (hydrated aluminium oxide (Al_2 O_3∙2H_2 O). The mixture is agitated and a stream of air is blown in which a floating froth. This is then filtered off leaving undisolved impurities.
Roasting – the metal ore is roasted in air to convert the sulphates and carbonate into oxides which can then be reduced to the metal.
Method of extraction
Summary
Extraction of metals
Metal Main occurrence Main method of extraction Equation of extraction
Ca,Ca^(2+))
Lime stone
CaCO_3
Gypsum
CaSO_4 Electrolysis of fused CaCl_2 and CaF_2 M^(n+)+ne ̅ → M
Na,Na^+)
Rock salt
NaCl
Chile slat petre
NaNO_3 Electrolysis of fused (molten), NaCl with CaCl_2 added M^(n+)+ne ̅ → M
Al Bauxite
Al_2 O_3∙2H_2 O
Silicate rocks Electrolysis of Al_2 O_3 in molten
Na_3 AlF_6
Zn Zinc blende
ZnS
Calamine
ZnCO_3 Reaction of zinc C or electrolysis of ZnSO_4 ZnOH+C →Zn+CO
Fe Magnetite
Fe_2 O_4
Haematite
Fe_2 O_3 Reduction of oxide with CO Fe_2 O_3+ 3CO →2Fe +3CO_2
Pb Galena PbS Reduction of PbO with C PbO+C → Pb¬+CO
Cu Copper pyrites
CuFeS
Cuprite Cu_2 O Partial oxidation of sulphide ore 2Cu_2 O + Cu_2 S →6Cu +SO_2
Hg Cinnabar
HgS Direct reduction of HgS by heat HgS+O_2→Hg+SO_2
Sodium
Sodium may be extracted from rock salt (sodium chloride) from seas and lakes.
Sodium nitrate (NaNO_3) and a double salt (NaHCO_2∙Na_2 CO_3.2H_2 O).
Extraction of sodium
Sodium is extracted by the down process using the down’s cell.
The cell consist of a carbon (graphite) anode and a circular (steel) cathode.
The anode is shielded from the cathode by chlorine from the cathode by an iron gauze to prevent sodium and chlorine from recombining after separation. The melting point of molten rock salt is about 8000C. this is lowed to about 6000C by adding calcium chloride to the molten salt.
Down’s cell for extraction of sodium
During electrolysis, sodium dissociate into sodium and chloride ions. Sodium ions, Na^+, (move to the cathode where they are reduced (accept electrons) to sodium atoms.
At anode
Na_((l)) + → Na_((l))
Sodium metal is then removed through an inverted trough placed over the anode.
Chlorine gas which is produced as a bi product is collected at the cathode.
At cathode
Cl_((aq))^- → Cl_((g)) + e
The chlorine is collected and taken away so that it does not recombine with the sodium formed at the anode.
Properties of sodium
Melting point Boiling point Density Electrical conductivity
97.80C 883 0.97 2.18 X 10-7C
Chemical properties
Tarnishes readily when exposed to air forming sodium oxide. The oxide absorbs moisture and form sodium hydroxide.
4Na_((s)) + O_(2(g)) → 2Na_2 O_((s))
Na_2 O_((s)) + H_2 O_((l)) → 2NaOH_((aq))
The sodium hydroxide is deliquescent hence absorbs carbondioxide to form sodium carbonate.
2NaOH_((aq)) + CO_(2(g)) → Na_2 CO_(3(s)) + H_2 O_((l))
Sodium burns in air with a golden yellow flame forming sodium peroxide.
2Na_((s)) + O_(2(g)) → Na_2 O_(2(s))
Sodium reacts violently with water forming an alkali and hydrogen gas is evolved.
2Na_((s)) + 2H_2 O_((l)) → 2NaOH_((aq)) + H_(2(g))
The reaction of sodium with steam is very explosive.
Sodium reacts with chlorine forming white fumes of sodium chloride.
2Na_((s)) + Cl_(2(g)) → 2NaCl
It reacts with ammonia to form sodamide
2Na_((s)) + 2NH_(3(g)) → 2NaNH_(2(s)) + H_(2(g))
Reacts explosively with acids forming salts
Reacts with mercury, forming sodium amalgam.
Na_((S)) + Hg_((S)) → Na/Hg_((s))
Uses of sodium
Sodium is used in industry in the following ways:
As an alloy with lead in the preparation of tetra ethyl lead(IV), which is an anti knock agent in petrol engines. This makes engines run more smoothly.
In the preparation of sodamide and sodium cynide (NaCN) used in the isolation of gold from metallic ores. Sodamide is used in the organic synthesis reactions such as the production of a dye called indigo.
To prepare sodium peroxide used to prepare hydrogen peroxide.
Liquid sodium is used as a heat conductor in nuclear reactors. It acts as a coolant because it is a good conductor of heat.
As a vapour in sodium lamps for street lights.
In making allows
In the production of titanium
In the production of sodium hydroxide used for see sinclani, B163
Extraction of iron; Fe
Iron ores are
Haematiti (Fe_2 O_3) or iron(III) oxide;
Magnetite (magnetic iron) Fe_3 O_4 (triontetraoxide)
Iron pyrites (iron disulphide) FeS_2
Siderite (Iron(II) carbonate) or spathetic iron FeCO_3
Limonite FeO(OH)
Haematite and magnetite are the most often used ores for the extraction of iron.
Raw materials and their uses
Coke (carbon) – burns to form carbon monoxide which is a reducing agnet.
Limestone (CaCO_(3))- forms calcium oxide which reacts with silica to form slag.
Iron ore (Fe_2 O_3,Fe_3 O_4 or FeCO_3)
Air / Conditions: Temperature at bottom of furnace, 17000C at top of furnace 7000C
Blas furnace
Iron is extracted in the blast furnace and the following processes and reactions take place.
The Fe_3 O_4 is mixed with coke (carbon and limestron (CaCO_3), this mixture is called cahrge, and introdued into the blast furnace. A blast of hot air can be introduced low down in the furnace through several pipes known as layers.
Reactions which take place in the furnace.
Stage 1
Formation of carbon dioxide
As the hot air comes in to contact with white-hot coke, the coke burns to form carbondioxide.
C_2 + O_((g)) → CO_(2(g)) DH= -39515Kmol
The above reaction liberates very large quantities of heat and it is this heat which keeps up the high temperature necessary for the production process (about 17000C)
Stage 2
Formation of carbondioxide
The carbondioxide formed above rises up the furnace where it encounters more coke. It is reduced to carbon monoxide.
CO_(2(g)) + C_2 → 2CO_((g))
Stage 3
Reduction of iron(III) oxide
The carbon monoxide reduces iron oxide to molten iron. The reactions are reversible and excess carbon monoxide is necessary to reduce the oxides completely.
Fe_2 O_3 + 3CO ⇌ 2Fe_((l)) + 3CO_(2(g)) (if haematiti is used)
Or
Fe_3 O_3 + 4CO ⇌ 3Fe_((l)) + 4CO_(2(g)) (If magnate is used)
Note the equations carefully
Functions of limestone CaCO_3)
Decomposition of limestone
The limestone CaCO_3) decomposes into calcium oxide (CaU) and carbon dioxide
CaCO_(3(s)) → CaO_((s)) + CO_(2(g))
The calcium oxide combines with sand (silicon (IV) oxide, SiO_2) present as an impurity in the ore and forms slag, calcium silicate.
CaO_((s)) + SiO_(2(s)) → CaSiO_(2(l))
Calcium silicate (slag)
Note: Both slag and irons are molten and drop to the bottom of the furnace. The less dense slag floats on top of iron and prevent the oxidation of iron by the air blast. The slag and iron are run out of separate holes.
The carbondioxide formed upon reduction of iron above is recycle and reacts with more coke © to form more carbondioxide.
Summary of reactions
Coke burns in the blast: C_((s)) + 〖 O〗_2 → CO_(2(g))
Limestone decomposes: CaCO_(3(s)) → CaO + CO_2
The carbon dioxide produced react with more coke: CO_(2(g)) + C_((s)) → 2CO_((g))
The iron ore is reduced by carbon monoxide gas:
Fe_2 O_(3(s)) + 3CO_((g)) → 2Fe_((l)) + 3CO_(2(g))
Earthly impurities in the ore (mainly silice) react with calcium oxide to form slag:
CaO + SiO_2 → CaSiO_2
slag
The blast furnace
Types and uses of iron
Pig iron – the iron produced in the blast furnace is about 90 – 95 percent pure. It is hard and brittle because it contains impurities like carbon and phosphorous, impurities lower its melting point to about 12000C; pure iron melts at 15390C.
Pig iron can be cast into various shapes by pouring molten metal into moulds. It makes clear castings because it expands on solidification, and it is called cast iron and has a flow tensile strength. It is used to make cookers, stoves, hot water radiators, railings, water pipes, posts, bases of bansens and other metal parts that need not be strong. It cannot be used for articles that require strength. e.g. bridges, metal bars for storage builds, etc.
Wrought iron
This is the purest form of commercial iron. It is 99% Fe_2 O_3, in a furnace. The oxides oxidize most of the impurities to gaseous oxides (CO_2,SO_2,P_2 O_5 ) and others to slag e.g. calcium silicate, which is removed by pressing the hot metal.
Fe_2 O_(3(s)) + 3C_((s)) → 2Fe_((s)) + 3CO_((g))
Uses
Wrought is strong and malleable (not brittle, like cast iron). Can be shaped by hammering at 10000C. It is used for making nails, sheets, chains, horse shoes, gates, farm machines and in cores of electromagnet because it cannot be permanently magnetized.
Steel
The major use of iron is its conversion to steel. Steel is an alloy of pure iron with carbon and chromium and tungsten. There are many types of steel with various uses. These are many types of steel with various uses. These can be summarized in the table below.
Type of steel Property Composition Use
Manganese steel Very tough Contains up to 13% manganese Rock breaking machines, rail cross overs
Stainless steel Does not rust Contains 20% chromium and 10% nickel Cutlery, surgical instruments, car bumpers
High speed steel Very tough and less brittle, metal parts can be machined up to 20 times faster with steel Iron 75.7%
Tungsten 18%
Chromium 6%
Vanadium 0.3% Cutting edges with lathes for metal work
Steel can also be used for building and construction industry for multi-storage buildings, bridges, farm equipment and a host of other uses.
Copper (Cu)
Occurrence
The main ores of copper are
Copper pyrites (CuFeS_2 ) or chalcopyrite
Malachite CuCO_3∙Cu(OH)_2
Cuprite Cu_2 O
Azurite 2CuCO_3∙CU(OH)
Copper glance (Cu_2 O)
Extraction
Copper pyrites CuFeS_2, is the ore usually used for the extraction f the metal. Three stages are involved:
The concentration of the ore to remove impurities
Roasting of the ore and reduction
Refining of the impure copper.
Copper pyrites is crushed to a powder and mixed with mixture and the heavier impurities sink to the bottom of the container. The froth on the surface, containing the concentrated ore, is separated off. The concentrate is then roasted in air in a furnace, producing copper(I) sulphide, iron(II) oxide and sulphur dioxide gas.
Equation for the reaction
2CuFeS_2 → Cu_2 S_((l)) + 2FeO_((l)) + 3SO_(2(g))
FeO_((l)) + SiO_(2(s)) → FeSiO_(3(s))
The sulphur oxide gas escapes from the top of the furnace. Iron(II) is removed by heating it in the absence of air with silica which is added in the furnace; iron(II) silicate is formed as 5kg which floats on top of molten copper(I) sulphide.
The copper(I) sulphide is finally reduced to the metal by hearing in a controlled supply of air. The impure copper formed is called blister copper because of the appearance like blisters on the copper surface on cooling caused by escaping gases.
Cu_2 S_((l)) + O_(2(g)) → 2Cu_((l)) + SO_(2(g))
Purification of blister copper
The impure copper is purified by electrolysis, impure copper bars form the anode and pure copper metal form the cathode. The electrolyte is a solution of copper(II) sulphate
At the anode
Pure copper from the impure anode goes into solution, reducing the size of anode.
Cu_((S)) → Cu_((aq))^(2+) + 2e
At the cathode
Copper ins accept two electrons each and are deposited as pure copper metal on the copper electrode which increase in size.
Cu_((aq))^(2+) + 2e → Cu_((s))
The purification of copper
Uses of copper
Pure copper
Copper is a good conductor of electricity and heat and so used in making electrical wires and cables, boilers, kettles, electric motors and dynamos.
Alloys of copper
Mixed with other metal copper can make alloys. E.g. brass and bronze
Brass is an alloy of copper (60%) and zinc (40%)
Brass is used for making gears, ships, propellers and rudders, bearings and statues.
Brass is also used to make screws, nuts, bolts, tubes, rods and ornaments.
Bronze – contains 90% copper and 10% tin. Other elements can be added for various reasons. It can be used for making ornaments.
Bronze can also be used for making gears, ships, propellers and rudders, bearings and statues.
Copper(II) sulphate is a constituent of many fungicides and presertives.
ALLOYS
An alloy is metallic substance consisting of a mixture of two or more metals or of a metal and non-metal. Allows have properties that are different from their components and make them, therefore more suitable for particular purposes than the parent metals.
Examples of commonly alloys are
Alloy Composition Uses
Brass Copper (60%), zinc (40%) For making gears, jewellery, ships’ propellers, rudders, brass band instruments, bearings and statues
Duralium or magnalium Copper (90%) with magnesium, manganese, silicon and copper Srews, casings, nuts, bolts, tubes, rods, and ornaments
Soft solder Tin (30%) and lead (20%) Connecting electrical wiring and other metal parts
Type metal Lead (60%) antimony (30%) and tin (10%) Casting, machine parts
Steel Iron with carbon and other elements See use of steel
REVISION QUESTIONS
TYPICAL UNEB PAST PAPERS
(a) “Extraction of metals is essentially a reduction process.” Explain the statement using extraction of iron as an example. Write an equation to illustrate your answer.
(b) State the conditions under which iron may react with:
(i) Oxygen
(ii) Water
(iii) Chlorine
(c) Write an equation for each of the reactions in (b) (ii) and (iii)
(d) Steel is an alloy of iron
(i) Explain what is meant by an alloy
(ii) Name the elements that are used to make stainless steel.
(iii) State one use of stainless steel
(iv) Suggest a reason why the use of stainless steel is preferred to that of pure iron. (Qn. 12, UNEB, 1999)
The diagram below shows an arrangement of the apparatus for the purification of copper.
Name the substance used as
Anode
Cathode
Name the electrolyte
Write equation for the reaction that took place at
Anode
Cathode (Qn. 2, UNEB, 2000)
(a) (i) What is an alloy?
(ii) Give an example of
State the composition of the alloy you have named in (a) (i)
(b) State two uses of the alloy (b) (ii) (Qn. 4, UNEB, 2000)
(a) Name one ore of iron and write its formula.
(b) During the extraction of iron, limestone and coke are added into the blast furnace. Explain the role of
(i) coke
(ii) limestone
(use equations to illustrate your answer)
(c) Write an equation leading to the formation of iron(II) sulphate
(d) Iron(II) sulphate was heated strongly
(i) State what was observed
(ii) Write equation for the reaction (Qn. 14, UNEB, 2002)
Duralumin is an alloy of aluminium, copper and element D.
(i) Identify element D
State one use of duralumin
Name the elements commonly used for making each of the following alloys and in each case, give one use of the alloy.
Steel
Elements
Use
Solder
Elements
Use
State tow reasons why alloys are commonly used instead of pure elements (Qn. 1, UNEB, 2015)
During manufacture of sodium hydroxide, concentrated sodium chloride (brine) solution is electrolyzed using mercury as the cathode.
(a) (i) Name the substance that is used as the anode
(ii) Give a reason for the choice of the substance
(iii) Identify the product collected at the cathode
(b) During the electrolysis, sodium amalgam is formed at the cathode
(i) State how sodium amalgam is converted to sodium hydroxide.
(ii) Write an equation leading to the formation of sodium hydroxide.
(c) State one industrial use of sodium hydroxide (Qn. 8, UNEB, 2015)
(a) Name the raw materials which are used in the extraction of iron using a blast furnace.
(b) Briefly describe the reactions that lead to the formation of iron during the extraction using a blast furnace. (Your answer should include equations for the reactions)
(c) State what would be observed and write equations for the reactions that would take place when the following gases are passed over heated iron
(i) Dry chlorine
(ii) Steam
Dilute hydrochloric acid was added to iron fillings and the mixture warmed. Write the equation for the equation that took place.
(Qn. 11, UNEB, 2012)
