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CHEMISTRY OF NITROGEN

A.NITROGEN

a) Occurrence:

Nitrogen is found in the atmosphere occupying about 78% by volume of air.

Proteins, amino acids, polypeptides in living things contain nitrogen.

b) Isolation of nitrogen from the air.

Nitrogen can be isolated from other gases present in air like oxygen, water (vapour), carbon (IV) oxide and noble gases as in the school laboratory as in the flow chart below:

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Water is added slowly into an “empty flask” which forces the air out into another flask containing concentrated sulphuric (VI) acid. Concentrated sulphuric (VI) acid is hygroscopic. It therefore absorb/remove water present in the air sample.

More water forces the air into the flask containing either concentrated sodium hydroxide or potassium hydroxide solution. These alkalis react with carbon IV) oxide to form the carbonates and thus absorbs/remove carbon IV) oxide present in the air sample.

 

Chemical equation 2NaOH (aq)  + CO2 (g)  ->  Na2CO3 (aq)  +  H2O(l)

Chemical equation 2KOH (aq)  + CO2 (g)  ->  K2CO3 (aq)  +  H2O(l)

More water forces the air through a glass tube packed with copper turnings. Heated brown copper turnings react with oxygen to form black copper (II) oxide.

 Chemical equation 2Cu (s)  +  O2 (g)  ->  CuO (s)  

(brown) (black)

The remaining gas mixture is collected by upward delivery/downward displacement of water/over water. It contains about 99% nitrogen and 1% noble gases.

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On a large scale for industrial purposes, nitrogen is got from fractional distillation of air.

c) Nitrogen from fractional distillation of air.

For commercial purposes nitrogen is got from the fractional of air.

Air is first passed through a dust precipitator/filter to remove dust particles.

The air is then bubbled through either concentrated sodium hydroxide or potassium hydroxide solution to remove/absorb Carbon(IV) oxide gas.

Chemical equation 2NaOH (aq) + CO2 (g)  ->  Na2CO3 (aq)  + H2O(l)

Chemical equation 2KOH (aq)  + CO2 (g)  ->  K2CO3 (aq)  + H2O(l)

Air mixture is the cooled to -25oC.At this temperature, water (vapour ) liquidifies and then solidify to ice and thus removed.

The air is further cooled to -200oC during which it forms a blue liquid.

The liquid is then heated. Nitrogen with a boiling point of -196oC distils first then Argon at-186oC and then finally Oxygen at -183oC boils last.

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c) School laboratory preparation of Nitrogen.

The diagram below shows the set up of the school laboratory preparation of nitrogen gas.

 

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d.Properties of Nitrogen gas(Questions)

1.Write the equation for the reaction for the school laboratory preparation of nitrogen gas.

Chemical equation NH4Cl (s) + NaNO2(s)->NaCl (g)+ NH4NO2 (s)

Chemical equation NH4NO2 (s)  ->  N2 (g)  +  H2O (l)

 2. State three physical properties of nitrogen gas.

colourless, odourless, less dense than air ,neutral and slightly soluble in water

3. State and explain the observation made when a burning magnesium ribbon is lowered in a gas jar containing nitrogen gas.

Observation; It continues burning with a blight blindening flame forming white ash.

Explanation

Magnesium burns to produce enough heat /energy to reacts with nitrogen to form white magnesium nitride.

 Chemical equation3Mg (s)  +  N2 (g)  ->  Mg3N2 (s)  

  (white ash/solid)

4. State two main uses of nitrogen gas

-manufacture of ammonia from Haber process

– As a refrigerant in storage of semen for Artificial insemination.  

B. OXIDES OF NITROGEN

Nitrogen forms three main oxides:

 i)Nitrogen(I) oxide(N2O)

 ii) Nitrogen(II) oxide (NO)

 iii) Nitrogen (IV) oxide( NO2)  

i) Nitrogen (I) oxide(N2O)

a) Occurrence

Nitrogen (I) oxide does not occur naturally but prepared in a laboratory.

b)Preparation

The set up below shows the set up of apparatus that can be used to prepare Nitrogen (I) oxide in a school laboratory.

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c) Properties of nitrogen (I) oxide (Questions)

1. Write the equation for the reaction for the school laboratory preparation of Nitrogen (I) oxide.

Chemical equation  NH4NO2(s)  ->  H2O (l)  + N2O (g)

2.a) State and explain three errors made in the above set up

 –Oxygen is being generated instead of Nitrogen (I) oxide.

 Ammonium Nitrate(V) should be used instead of potassium manganate(VI) and manganese(IV)oxide.

b) State three physical properties of Nitrogen (I) oxide.

-slightly soluble in water.

-colourless

-odourless

-less dense than air

-slightly sweet smell

3. State and explain the observation made when a burning magnesium ribbon is lowered in a gas jar containing Nitrogen (I) oxide.

Observation – Continues to burn with a bright flame

-White solid/residue is formed

Explanation-Magnesium burns in air to produce enough heat/energy split/break Nitrogen (I) oxide gas into free Nitrogen and oxygen then continues to burn in oxygen to form white solid/ash of Magnesium oxide.

Chemical equation

 Mg(s)  + N2O (g)->  MgO (s)  + N2(g)  

 

4. State and explain the observation made when the following non metals are burnt then lowered in a gas jar containing Nitrogen (I) oxide.

 a) Carbon/charcoal

Observation – Continues to burn with an orange glow

-colorless gas is formed that forms white precipitate with lime water.

Explanation-Carbon/charcoal burns in air to produce enough heat/energy split/break Nitrogen (I) oxide gas into free Nitrogen and oxygen then continues to burn in oxygen to form carbon (IV) oxide gas. Carbon (IV) oxide gas reacts to form a white precipitate with lime water.

Chemical equation C(s)  + 2N2O (g)->  CO2 (g)  + 2N2(g)  

 b) Sulphur powder

Observation – Continues to burn with a blue flame

-colorless gas is formed that turn orange acidified potassium dichromate (VI) to green.

Explanation-Sulphur burns in air to produce enough heat/energy split/break Nitrogen (I) oxide gas into free Nitrogen and oxygen then continues to burn in oxygen to form sulphur (IV) oxide gas. Sulphur (IV) oxide gas turns orange acidified potassium dichromate (VI) to green.

Chemical equation S(s)  + 2N2O (g)->  SO2 (g)  + 2N2(g)  

5. State two uses of nitrogen (I) oxide

 -As laughing gas because as anesthesia the patient regain consciousness laughing hysterically after surgery.

 -improves engine efficiency.

6. State three differences between nitrogen (I) oxide and oxygen

 –Oxygen is odourless while nitrogen (I) oxide has faint sweet smell

-Both relight/rekindle a glowing wooden splint but Oxygen can relight a feeble glowing splint while nitrogen (I) oxide relights well lit splint.

-Both are slightly soluble in water but nitrogen (I) oxide is more soluble.

ii) Nitrogen (II) oxide (NO)

a) Occurrence

Nitrogen (II) oxide does not occur naturally but prepared in a laboratory.

b)Preparation

The set up below shows the set up of apparatus that can be used to prepare Nitrogen (II) oxide in a school laboratory.

Image From EcoleBooks.com

  c) Properties of nitrogen (II) oxide (Questions)

 

  1. Write the equation for the reaction for the school laboratory preparation of Nitrogen (II) oxide.

Chemical equation 3Cu(s) + 8HNO3(aq) -> 4H2O (l)+2NO (g) +2Cu(NO3)2(aq)

Chemical equation 3Zn(s) + 8HNO3(aq) -> 4H2O (l)+2NO (g) +2Zn(NO3)2(aq)

Chemical equation 3Mg(s) + 8HNO3(aq) -> 4H2O (l)+2NO (g)+2Mg(NO3)2(aq)

2. State three physical properties of Nitrogen (II) oxide.

-insoluble in water.

-colourless

-odourless

-denser dense than air

-has no effect on both blue and red litmus papers

  1. State and explain the observation made when a burning magnesium ribbon is lowered in a gas jar containing Nitrogen (II) oxide.

Observation – Continues to burn with a bright flame

-White solid/residue is formed

Explanation-Magnesium burns in air to produce enough heat/energy split/break Nitrogen (II) oxide gas into free Nitrogen and oxygen then continues to burn in oxygen to form white solid/ash of Magnesium oxide.

Chemical equation 2Mg(s)  + 2NO (g)->  2MgO (s)  + N2(g)  

  1. State and explain the observation made when the following non metals are burnt then lowered in a gas jar containing Nitrogen (II) oxide.

 a) Carbon/charcoal

Observation – Continues to burn with an orange glow

-colorless gas is formed that forms white precipitate with lime water.

Explanation-Carbon/charcoal burns in air to produce enough heat/energy split/break Nitrogen (II) oxide gas into free Nitrogen and oxygen then continues to burn in oxygen to form carbon (IV) oxide gas.Carbon (IV) oxide gas reacts to form a white precipitate with lime water.

Chemical equation C(s)  + 2NO (g)->  CO2 (g)  + N2(g)  

 b) Sulphur powder

Observation – Continues to burn with a blue flame

-colorless gas is formed that turn orange acidified potassium dichromate (VI) to green.

Explanation-Sulphur burns in air to produce enough heat/energy split/break Nitrogen (II) oxide gas into free Nitrogen and oxygen then continues to burn in oxygen to form sulphur (IV) oxide gas.Sulphur (IV) oxide gas turns orange acidified potassium dichromate (VI) to green.

Chemical equation S(s)  + N2O (g)->  SO2 (g)  + N2(g)  

c) Phosphorus

Observation – Continues to produce dense white fumes  

Explanation-Phosphorus burns in air to produce enough heat/energy split/break Nitrogen (II) oxide gas into free Nitrogen and oxygen then continues to burn in oxygen to form dense white fumes of phosphorus (V) oxide gas.

Chemical equation 4P(s)  + 10NO (g)->  2P2O5(g)  + 5N2(g)  

5. State one use of nitrogen (II) oxide

As an intermediate gas in the Ostwalds process for manufacture of nitric(V) gas.

6. State and explain the observation made when nitrogen (II) oxide is exposed to the atmosphere.

Observation–brown fumes produced/evolved that turn blue litmus paper red.

Explanation- Nitrogen (II) oxide gas on exposure to air is quickly oxidized by the air/ oxygen to brown nitrogen (IV) oxide gas. Nitrogen (IV) oxide gas is an acidic gas.

Chemical equation 2NO (g)+ O2(g)->  2NO2 (g)

 (colorless)  (brown)

ii) Nitrogen (IV) oxide (NO2)

a) Occurrence

Nitrogen (IV) oxide occurs -naturally from active volcanic areas.

-formed from incomplete combustion of the internal combustion engine of motor vehicle exhaust fumes.

-from lightening

b)Preparation

The set up below shows the set up of apparatus that can be used to prepare Nitrogen (IV) oxide in a school laboratory.

 

Image From EcoleBooks.com

c) Properties of nitrogen (IV)oxide (Questions)

1. Write the equation for the reaction for the school laboratory preparation of Nitrogen (II) oxide.

Chemical equation Cu(s) + 4HNO3(aq) -> 2H2O (l)+2NO 2(g) +Cu(NO3)2(aq)

Chemical equation Zn(s) + 4HNO3(aq) -> 2H2O (l)+2NO 2(g) +Zn(NO3)2(aq)

Chemical equation Fe(s) + 4HNO3(aq) -> 2H2O (l)+2NO 2(g) +Fe(NO3)2(aq)

2. State three physical properties of Nitrogen (IV) oxide.

-soluble/dissolves in water.

-brown in colour

-has pungent irritating poisonous odour/smell

-denser dense than air

-turns blue litmus papers to red

3. State and explain the observation made when Nitrogen (IV) oxidegas is bubbled in water.

Observation–The gas dissolves and thus brown colour of the gas fades

-A colourless solution is formed

-solution formed turns blue litmus papers to red

 -solution formed has no effect on red

Explanation-Magnesium burns in air to produce enough heat/energy split/break Nitrogen (IV) oxide gas dissolves then react with water to form an acidic mixture of nitric(V) acid andnitric(III) acid.

Chemical equationH2O (l) + 2NO 2(g)->HNO3(aq) + HNO2(aq)

 (nitric(V) acid) (nitric(III) acid)

4. State and explain the observation made when a test tube containing Nitrogen (IV) oxide is cooled then heated gently then strongly.

Observation on cooling

 -Brown colour fades

 -Yellow liquid formed

Observation on gentle heating

  –Brown colour reappears

Yellow liquid formed changes to brown fumes/gas

Observation on gentle heating

  –Brown colour fades

brown fumes/gas changes to a colourless gas

Explanation-Brown nitrogen (IV) oxide gas easily liquefies to yellow dinitrogen tetraoxide liquid.When the yellow dinitrogen tetraoxide liquid is gently heated it changes back to the brown nitrogen (IV) oxidegas.When the brown nitrogen (IV) oxide gas is strongly heated it decomposes to colourless mixture of Nitrogen (II) oxide gas and Oxygen.

Chemical equation O2(s) + 2NO (g) ===== 2NO2 (g) ===== N2O4(l)  

(colourless gases) (brown gas) (yellow liquid)

5. State and explain the observation made when a burning magnesium ribbon is lowered in a gas jar containing Nitrogen (IV) oxide.

Observation – Continues to burn with a bright flame

-White solid/residue is formed

-Brown fumes/colour fades

Explanation-Magnesium burns in air to produce enough heat/energy split/break brown Nitrogen (IV) oxide gas into free colourless Nitrogen and oxygen then continues to burn in oxygen to form white solid/ash of Magnesium oxide.

Chemical equation 4Mg(s)  + 2NO 2(g)->  4MgO (s)  + N2(g)

4. State and explain the observation made when the following non metals are burnt then lowered in a gas jar containing Nitrogen (IV) oxide.

 a) Carbon/charcoal

Observation – Continues to burn with an orange glow

-Brown fumes/colour fades

-colorless gas is formed that forms white precipitate with lime water.

Explanation-Carbon/charcoal burns in air to produce enough heat/energy split/break brown Nitrogen (IV) oxide gas into free colourless Nitrogen and oxygen then continues to burn in oxygen to form carbon (IV) oxide gas.Carbon (IV) oxide gas reacts to form a white precipitate with lime water.

Chemical equation2C(s)  + 2NO 2(g)->  2CO2 (g)  + N2(g)  

 b) sulphur powder

Observation – Continues to burn with a blue flame

  -Brown fumes/colour fades  

-colorless gas is formed that turn orange acidified potassium dichromate (VI) to green.

Explanation-Sulphur burns in air to produce enough heat/energy split/break brown Nitrogen (IV) oxide gas into free colourless Nitrogen and oxygen then continues to burn in oxygen to form sulphur (IV) oxide gas.Sulphur (IV) oxide gas turns orange acidified potassium dichromate (VI) to green.

Chemical equation2S(s)  + 2NO2 (g)->  2SO2 (g)  + N2(g)  

c) Phosphorus

Observation- Continues to produce dense white fumes

-Brown fumes/colour fades

Explanation-Phosphorus burns in air to produce enough heat/energy split/break brown Nitrogen (IV) oxide gas into free colourless Nitrogen and oxygen then continues to burn in oxygen to form dense white fumes of phosphorus (V) oxide gas.

Chemical equation 8P(s)  + 10NO2 (g)->  4P2O5(g)  + 5N2(g)  

5. State two uses of nitrogen (IV) oxide

  -In theOstwald process for industrial manufacture of nitric (V) gas.

  -In the manufacture of T.N.T explosives

6. State and explain the observation made when nitrogen (II) oxide is exposed to the atmosphere.

Observation–brown fumes produced/evolved that turn blue litmus paper red.

Explanation- Nitrogen (II) oxide gas on exposure to air is quickly oxidized by the air/ oxygen to brown nitrogen (IV) oxide gas. Nitrogen (IV) oxide gas is an acidic gas.

Chemical equation 2NO (g) + O2(g) ->   2NO2 (g)

 (colourless)  (brown)

 
 

C. AMMONIA (NH3)

Ammonia is a compound of nitrogen and hydrogen only. It is therefore a hydride of nitrogen.

a) Occurrence

Ammonia gas occurs -naturally from urine of mammals and excretion of birds

-formed in the kidney of human beings

b)Preparation

The set up below shows the set up of apparatus that can be used to prepare dry Ammonia gas in a school laboratory.

Set up method 1

Image From EcoleBooks.com

 

1. Write the equation for the reaction taking place in:

  1. Method 1

Chemical equation

Ca (OH)2(s)  + NH4 Cl(s)->CaCl2 (aq) + H2O(l)  + 2NH3(g)

 b)Method 2

Chemical equation

NaOH (aq)  + NH4 Cl(aq) -> NaCl (aq) + H2O(l)  + NH3(g)

2. State three physical properties of ammonia.

 -has a pungent choking smell of urine

 -Colourless

 -Less dense than air hence collected by upward delivery

-Turns blue litmus paper blue thus is the only naturally occurring basic gas (at this level)

3. Calcium oxide is used as the drying agent. Explain why calcium chloride and concentrated sulphuric(VI) acid cannot be used to dry the gas.

-Calcium chloride reacts with ammonia forming the complex compound CaCl2.8H2O.

Chemical equation CaCl2 (s)  + 8NH3(g) -> CaCl2 .8NH3(g)

-Concentrated sulphuric(VI) acid reacts with ammonia forming ammonium sulphate(VI) salt compound

 Chemical equation 2NH3(g)  +H2SO4(l) ->(NH4)2SO4(aq)

4. Describe the test for the presence of ammonia gas.

Using litmus paper:

Dip moist/damp/wet blue and red litmus papers in a gas jar containing a gas suspected to be ammonia.The blue litmus paper remain blue and the red litmus paper turns blue.Ammonia is the only basic gas.(At this level)

Using hydrogen chloride gas

Dip a glass rod in concentrated hydrochloric acid. Bring the glass rod near the mouth of a gas jar suspected to be ammonia. White fumes (of ammonium chloride)are produced/evolved.

5. Describe the fountain experiment to show the solubility of ammonia.

Ammonia is very soluble in water.

When a drop of water is introduced into flask containing ammonia, it dissolves all the ammonia in the flask. If water is subsequently allowed into the flask through a small inlet, atmospheric pressure forces it very fast to occupy the vacuum forming a fountain. If the water contains three/few drops of litmus solution, the litmus solution turns blue because ammonia is an alkaline/basic gas. If the water contains three/few drops of phenolphthalein indicator, the indicator turns pink because ammonia is an alkaline/basic gas. Sulphur(IV)oxide and hydrogen chloride gas are also capable of the fountain experiment . If the water contains three/few drops of phenolphthalein indicator, the indicator turns colourless because both Sulphur(IV) oxide and hydrogen chloride gas are acidic gases.

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6. State and explain the observation made when hot platinum /nichrome wire is placed over concentrated ammonia solution with Oxygen gas bubbled into the mixture.

Observations

Hot platinum /nichrome wire continues to glow red hot.

Brown fumes of a gas are produced.

Explanation

Ammonia reacts with Oxygen on the surface of the wire .This reaction is exothermic producing a lot of heat/energy that enables platinum wire to glow red hot. Ammonia is oxidized to Nitrogen(II)oxide gas and water. Hot platinum /nichrome wire acts as catalyst to speed up the reaction. Nitrogen(II)oxide gas is further oxidized to brown Nitrogen(IV)oxide gas on exposure to air.

Chemical equation

(i)4NH3(g)  +  5O2(g) -Pt-> 4NO(g) + 6H2O(l)

(ii)2NO(g) + O2(g) -> 2NO2(g)

7. Ammonia gas was ignited in air enriched with Oxygen gas. State and explain the observations made

Observations

 – Ammonia gas burns with a green flame

 -Colourless gas produced

Explanation

Ammonia gas burns with a green flame in air enriched with Oxygen to from Nitrogen gas and water.

Chemical equation

 2NH3(g)  +  O2(g) -> N2(g) + 3H2O(l)

8. Dry ammonia was passed through heated copper(II)Oxide as in the set up below.

Image From EcoleBooks.com

(a)State the observations made in tube K

-Colour changes from black to brown

-Colourless liquid droplet form on the cooler parts of tube K

(b)(i)Identify liquid L.

 -Water/ H2O(l)


(ii)Explain a chemical and physical test that can be used to identify liquid L.

 Chemical test

(i) Add three/few drops of liquid L into anhydrous copper(II)sulphate(VI).

  Colour changes from white to blue.

Explanation-Water changes white anhydrous copper(II)sulphate(VI) to blue hydrated copper(II)sulphate(VI)

  (ii) Add three/few drops of liquid L into anhydrous cobalt(II)Chloride.

  Colour changes from blue to pink.

Explanation-Water changes blue anhydrous cobalt(II)Chloride to pink hydrated cobalt(II)Chloride.

 Physical test

(i)Heat the liquid. It boils at 100oC at sea level (1atmosphere pressure/760mmHg pressure, 101300Pa,101300Nm-2).

 (ii)Cool the liquid. It freezes at 0.0oC .

 (iii)Determine the density. It is 1.0gcm-3


(c)Write the equation for the reaction that take place.

2NH3(g)  +  3CuO(s) -> N2(g) + 3H2O(l) + 3Cu(s)

  (black)  (brown)

2NH3(g)  +  3PbO(s) -> N2(g) + 3H2O(l) + 3Pb(s)

 (brown when hot) (grey)

8.(a)What is aqueous ammonia


Aqueous ammonia is formed when ammonia gas is dissolved in water.

NH3(g) + (aq)  -> NH3(aq)

A little NH3(aq) reacts with ammonia water to form ammonia solution(NH4OH)

NH3 (aq) + H2O(l) OH(aq) + NH4+(aq)

This makes a solution of aqueous ammonia is a weak base /alkali unlike other two alkalis.

9.Using dot and cross to represent outer electrons show the bonding in:

(a) NH3

 


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(b) NH4+

 

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(c)NH4Cl

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10.Name four uses of ammonia

 (i)In the manufacture of nitrogenous fertilizers.

 (ii) In the manufacture of nitric(V)acid from Ostwalds process.

 (iii)As a refrigerant in ships and warehouses.

 (iv)In softening hard water.

 (v)In the solvay process for the manufacture of sodium carbonate.

 (vi)In the removal of grease and stains.

11.(a)Calculate the percentage of Nitrogen in the following fertilizers:

 (i) (NH4)2SO4

Molar mass of (NH4)2SO4 = 132g

Mass of N in (NH4)2SO4= 28g

% of N => 28 x 100 = 21.2121%

 132

 (ii) (NH4)3PO4

   Molar mass of (NH4)3PO4 = 149g

Mass of N in (NH4)3PO4= 42g

% of N => 42 x 100 = 28.1879%

  149

(b)State two advantages of fertilizer a (i) over a (ii) above.

(i)Has higher % of Nitrogen

(ii)Has phosphorus which is necessary for plant growth.

(c) Calculate the mass of Nitrogen in a 50kg bag of:

 (i) (NH4)2SO4

% of N in (NH4)2SO4 = 21.2121%

Mass of N in 50 kg (NH4)2SO4= 21.2121 x 50 = 10.6 kg

  100

  (ii) NH4NO3

Molar mass of NH4NO3 = 80g

Mass of N in (NH4)3PO4= 28g

% of N => 28 x 100 = 35%

  80

% of N in NH4NO3 = 35%

Mass of N in 50 kg (NH4)2SO4= 35 x 50 = 17.5 kg

  100


NH4NO3 therefore has a higher mass of Nitrogen than (NH4)2SO4

d).Manufacture of Ammonia /Haber process

Most of the Ammonia produced for industrial purposes uses the Haber process developed by the German Scientist Fitz Haber.

(i)Raw materials

The raw materials include:

 (i)Nitrogen from Fractional distillation of air from the atmosphere.

 (ii)Hydrogen from:

I. Water gas-passing steam through heated charcoal

C(s) + H2O(l) -> CO(g) + H2 (g)

II .Passing natural gas /methane through steam.

CH4(g)+ H2O(l) -> CO(g) + 3H2 (g)

 

 

(ii)Chemical process

Hydrogen and Nitrogen are passed through a purifier to remove unwanted gases like Carbon(IV)oxide,Oxygen,sulphur(IV)oxide, dust, smoke which would poison the catalyst.

Hydrogen and Nitrogen are then mixed in the ratio of 3:1 respectively. The mixture is compressed to 200-250atmoshere pressure to liquidify. The liquid mixture is then heated to 400- 450oC.The hot compressed gases are then passed over finely divided Iron catalyst promoted/impregnated with Al2O3 /K2O .Promoters increase the efficiency of the catalyst.

Image From EcoleBooks.com

 

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Optimum conditions in Haber processs

Chemical equation

N2 (g) + 3H2 (g) ===Fe/Pt=== 2NH3 (g) ΔH = -92kJ

Equilibrium/Reaction rate considerations

(i)Removing ammonia gas once formed shift the equilibrium forward to the right to replace the ammonia. More/higher yield of ammonia is attained.

(ii)Increase in pressure shift the equilibrium forward to the right where there is less volume/molecules. More/higher yield of ammonia is attained. Very high pressures raise the cost of production because they are expensive to produce and maintain. An optimum pressure of about 200atmospheres is normally used.

(iii)Increase in temperature shift the equilibrium backward to the left because the reaction is exothermic (ΔH = -92kJ) . Ammonia formed decomposes back to Nitrogen and Hydrogen to remove excess heat therefore a less yield of ammonia is attained. Very low temperature decreases the collision frequency of Nitrogen and Hydrogen and thus the rate of reaction too slow and uneconomical.

An optimum temperature of about 450oC is normally used.

(iv)Iron and platinum can be used as catalyst. Platinum is a better catalyst but more expensive and easily poisoned by impurities than Iron. Iron is promoted /impregnated with AluminiumOxide(Al2O3) to increase its surface area/area of contact with reactants and thus efficiency. The catalyst does not increase the yield of ammonia but it speed up its rate of formation.

e) Nitric(V)acid (HNO3)

a)Introduction.

Nitric(V)acid is one of the mineral acids .There are three mineral acids; Nitric(V)acid, sulphuric(VI)acid and hydrochloric acid. Mineral acids do not occur naturally but are prepared in a school laboratory and manufactured at industrial level.

b) School laboratory preparation

Nitric(V)acid is prepared in a school laboratory from the reaction of Concentrated sulphuric(VI)acid and potassium nitrate(V) below.  Image From EcoleBooks.com

(c)Properties of Concentrated Nitric (V)acid(Questions)

1.Write an equation for the school laboratory preparation of nitric(V)acid.

 KNO3(s) + H2SO4(l)  -> KHSO4(s) + HNO3(l)

2.Sodium nitrate(V)can also be used to prepare nitric(V)acid. State two reasons why potassium nitrate(V) is preferred over Sodium nitrate(V).

(i) Potassium nitrate(V) is more volatile than sodium nitrate(V) and therefore readily displaced from the less volatile concentrated sulphuric(VI)acid

(ii)
Sodium nitrate(V) is hygroscopic and thus absorb water . Concentrated sulphuric(VI)acid dissolves in water. The dissolution is a highly exothermic process.

3. An all glass apparatus /retort is used during the preparation of nitric(V) acid. Explain.

Hot
concentrated nitric(V) acid vapour is highly corrosive and attacks rubber cork apparatus if used.

4. Concentrated nitric(V) acid is colourless . Explain why the prepared sample in the school laboratory appears yellow.

 

Hot concentrated nitric(V) acid decomposes to brown nitrogen(IV)oxide and Oxygen gases.

 4HNO3(l/g) -> 4NO2(g) + H2O (l) +O2(g)

Once formed the brown nitrogen(IV)oxide dissolves in the acid forming a yellow solution .

5. State and explain the observation made when concentrated nitric (V) acid is heated.

Observation

Brown fumes are produced.

Colourless gas that relights/rekindles glowing splint

Explanation

Hot concentrated nitric(V) acid decomposes to water, brown nitrogen(IV)oxide and Oxygen gases. Oxygen gas is not visible in the brown fumes of nitrogen (IV) oxide.

 4HNO3(g) -> 4NO2(g) + H2O (l) +O2(g)

6. Explain the observations made when:

(a) About 2cm3 of Iron(II)sulphate(VI) solution is added about 5 drops of concentrated nitric(V) acid and the mixture then heated/warmed in a test tube.

Observation

 (i)Colour changes from green to brown.

 (ii)brown fumes /gas produced on the upper parts of the test tube.

Explanation

 Concentrated nitric(V) acid is a powerful/strong oxidizing agent. It oxidizes green Fe2+ ions in FeSO4 to brown/yellow Fe3+ .The acid is reduced to colourless Nitrogen(II)oxide.

Chemical equation:

6FeSO4(aq) + 3H2SO4 (aq) + 2HNO3(aq) -> 3Fe2(SO4) 3 (aq)+ 4H2O + 2NO(g)

Colourless Nitrogen(II)oxide is rapidly further oxidized to brown Nitrogen(IV)oxide by atmospheric oxygen.

Chemical equation:

 2NO(g)  +   O(g)   ->  2NO2 (g)

 (colourless) (brown)


(b) A spatula full of sulphur powder in a clean dry beaker was added to 10cm3 concentrated nitric (V) acid and then heated gently/warmed.

Observation

 (i)Yellow colour of sulphur fades.

 (ii) Brown fumes /gas produced.

Explanation

 Concentrated nitric(V) acid is a powerful/strong oxidizing agent. It oxidizes yellow sulphur to colourless concentrated sulphuric(VI)acid. The acid is reduced to brown Nitrogen(IV)oxide gas.

Chemical equation:

 S(s) + 6HNO3(l) -> 4NO2(g) + H2O (l) +H2SO4(l)

(c) A few/about 1.0g pieces of copper turnings/Zinc granules/ Magnesium ribbon are added 10cm3 of concentrated nitric(V) acid in a beaker.


Observation

 (i) Brown fumes /gas produced.

 (ii) Blue solution formed with copper turnings  

 (iii) Colourless solution formed with Zinc granules/Magnesium ribbon  

Explanation

Concentrated nitric (V) acid is a powerful/strong oxidizing agent. It oxidizes metals to their metal nitrate (VI) salts. The acid is reduced to brown Nitrogen (IV) oxide gas.

Chemical equation:

 Cu(s) + 4HNO3(l) -> 2NO2(g) + H2O (l) + Cu(NO3) 2 (aq)

 Zn(s) + 4HNO3(l) -> 2NO2(g) + H2O (l) + Zn(NO3) 2 (aq)

 Mg(s) + 4HNO3(l) -> 2NO2(g) + H2O (l) + Mg(NO3) 2 (aq)

 Pb(s) + 4HNO3(l) -> 2NO2(g) + H2O (l) + Pb(NO3) 2 (aq)

 Ag(s) + 2HNO3(l) -> NO2(g) + H2O (l) + AgNO3 (aq)

(d)Properties of Dilute Nitric (V)acid(Questions)

(i)What is dilute nitric(v)acid

When concentrated nitric(v)acid is added to over half portion of water ,it is relatively said to be dilute. A dilute solution is one which has more solvent/water than solute/acid. The number of moles of the acid are present in a large amount/volume of the solvent.This makes the molarity /number of moles present in one cubic decimeter of the solution to be low e.g. 0.02M.

If more water is added to the acid until the acid is too dilute to be diluted further then an infinite dilute solution if formed.


(ii))1cm length of polished Magnesium ribbon was put is a test tube containing 0.2M dilute nitric(v)acid. State and explain the observation made.

Observation

-Effervescence/bubbling/fizzing

-Colourless gas produced that extinguish burning splint with an explosion/pop sound

-Colourless solution formed

-Magnesium ribbon dissolves/decrease in size

Explanation

Dilute dilute nitric(v)acid reacts with Magnesium to form hydrogen gas.

Mg(s) + 2HNO3(aq) -> H2 (g) + Mg(NO3) 2 (aq)

With other reactive heavy metals, the hydrogen gas produced is rapidly oxidized to water.

Chemical equation 3Pb(s) + 8HNO3(aq) -> 4H2O (l)+2NO (g) +2Pb(NO3)2(aq)

Chemical equation 3Zn(s) + 8HNO3(aq) -> 4H2O (l)+2NO (g) +2Zn(NO3)2(aq)

Chemical equation 3Fe(s) + 8HNO3(aq) -> 4H2O (l)+2NO (g) +2Fe(NO3)2(aq)

Hydrogen gas therefore is usually not prepared in a school laboratory using dilute nitric (v)acid.

(iii)A half spatula full of sodium hydrogen carbonate and Copper(II) carbonate were separately to separate test tubes containing 10cm3 of 0.2M dilute nitric (V) acid.

Observation

-Effervescence/bubbling/fizzing

-Colourless gas produced that forms a white precipitate with lime water.

-Colourless solution formed with
sodium hydrogen carbonate.

– Blue solution formed with
Copper(II) carbonate.

Explanation

Dilute dilute nitric (v)acid reacts with Carbonates and hydrogen carbonates to form Carbon(IV)oxide, water and nitrate(V)salt

CuCO3 (s) + 2HNO3(aq) -> H2O (l) + Cu(NO3) 2 (aq) + CO2 (g)

ZnCO3 (s) + 2HNO3(aq) -> H2O (l) + Zn(NO3) 2 (aq) + CO2 (g)

CaCO3 (s) + 2HNO3(aq) -> H2O (l) + Ca(NO3) 2 (aq) + CO2 (g)

PbCO3 (s) + 2HNO3(aq) -> H2O (l) + Pb(NO3) 2 (aq) + CO2 (g)

FeCO3 (s) + 2HNO3(aq) -> H2O (l) + Fe(NO3) 2 (aq) + CO2 (g)

NaHCO3 (s) + HNO3(aq) -> H2O (l) + NaNO3 (aq) + CO2 (g)

KHCO3 (s) + HNO3(aq) -> H2O (l) + KNO3 (aq) + CO2 (g)

NH4HCO3 (aq) + HNO3(aq) -> H2O (l) + NH4NO3 (aq) + CO2 (g)

Ca(HCO3) 2 (aq) + 2HNO3(aq) -> 2H2O (l) + Ca(NO3) 2 (aq) + 2CO2 (g)

Mg(HCO3) 2 (aq) + 2HNO3(aq) -> 2H2O (l) + Mg(NO3) 2 (aq) + 2CO2 (g)

(iii) 25.0cm3 of 0.1M Nitric(V) acid was titrated with excess 0.2M sodium hydroxide solution using phenolphthalein indicator.

 I. State the colour change at the end point

  Colourless

 II. What was the pH of the solution at the end point. Explain.

pH 1/2/3

  A little of the acid when added to the base changes the colour of the indicator to show the end point. The end point therefore is acidic with low pH of Nitric(V) acid. Nitric(V) acid is a strong acid with pH 1/2/3.

 III. Calculate the number of moles of acid used.

Number of moles = molarity x volume => 0.1 x 25 = 2.5 x 10-3moles

  1000 1000

 IV. Calculate the volume of sodium hydroxide used

Volume of sodium hydroxide in cm3

= 1000 x Number of moles =>   1000x 2.5 x 10-3 =  12.5cm3

 Molarity 0.2

(e)Industrial large scale manufacture of Nitric (V) acid

(i)Raw materials

1. Air/Oxygen

 Oxygen is got from fractional distillation of air

 Ammonia from Haber process.

2. Chemical processes

Air from the atmosphere is passes through electrostatic precipitators/filters to remove unwanted gases like Nitrogen, Carbon (IV) oxide, dust, smoke which may poison the catalyst. The ammonia -air mixture is compressed to 9 atmospheres to reduce the distance between reacting gases.

The mixture is passed through the heat exchangers where a temperature of 850oC-900oC is maintained.

The first reaction takes place in the catalytic chamber where Ammonia reacts with the air to form Nitrogen (II) Oxide and water.

Optimum condition in Ostwald’s process

Chemical equation

4NH3 (g) + 5O2 (g) === Pt/Rh === 4NO (g) + 6H2O (g) ΔH = 950kJ

The reaction is reversible and exists in dynamic equilibrium where the products reform back the reactants. The following factors are used to increase the yield/amount of Nitrogen (II) oxide:

(i)Removing Nitrogen (II) oxide gas once formed shift the equilibrium forward to the right to replace the Nitrogen (II) oxide.

More/higher yield of Nitrogen (II) oxide is attained as reactants try to return the equilibrium balance.

(ii)Increase in pressure shift the equilibrium backward to the left where there are less volume/molecules.

Less/lower yield of Nitrogen (II) oxide is attained.

Very low pressures increases the distance between reacting NH3 and O2 molecules.

An optimum pressure of about 9 atmospheres is normally used.

Cooling the mixture condenses the water vapour to liquid water

(iii)Increase in temperature shift the equilibrium backward to the left because the reaction is exothermic (ΔH = -950kJ).

Nitrogen (II) oxide and water vapour formed decomposes back to Ammonia and Oxygen to remove excess heat therefore a less yield of Nitrogen (II) oxide is attained.

Very low temperature decreases the collision frequency of Ammonia and Oxygen and thus the rate of reaction too slow and uneconomical.

An optimum temperature of about 900oC is normally used.

(iv)Platinum can be used as catalyst.  

Platinum is very expensive. It is:

-promoted with Rhodium to increase the surface area/area of contact.

-added/coated on the surface of asbestos to form Platonized –asbestos to reduce the amount/quantity used.

The catalyst does not increase the yield of Nitrogen (II) Oxide but it speed up its rate of formation.

Nitrogen (II) oxide formed is passed through an oxidation reaction chamber where more air oxidizes the Nitrogen (II) Oxide to Nitrogen (IV) Oxide gas.

Chemical equation

2NO (g)  +  O2 (g) -> 2NO2 (g)

Nitrogen (IV) Oxide gas is passed up to meet a downward flow of water in the absorption chamber. The gas reacts with water to form a mixture of Nitric (V) and Nitric (III) acids

Chemical equation.

2NO2 (g)  +  H2O (l) -> HNO2 (as) + HNO3 (as)

Excess air is bubbled through the mixture to oxidize Nitric (III)/ HNO2 (as) to Nitric (V)/HNO3 (as)


Chemical equation.

O2 (g)  +  2HNO2 (as) -> 2HNO3 (as)

Overall chemical equation in the absorption chamber.

O2 (g)  +  4NO2 (g) +  2H2O (l) -> 4HNO3 (as)

The acid is 65% concentrated. It is made 100% concentrated by either:

 (i) fractional distillation or

 (ii) added to concentrated sulphuric (VI) acid to remove the 35% of water.

Image From EcoleBooks.com

 

A factory uses 63.0 kg of 68% pure nitric (V) acid per day to produce an ammonium fertilizer for an agricultural county. If the density of the acid is 1.42 gcm-3, calculate:

(i) the concentration of the acid used in moles per litre.


Molar mass HNO3 = 63

Method 1

Moles of HNO3 in 1cm3 = Mass in 1cm3 1.42 => 1.42   = 0.0225 moles

  Molar mass HNO3 63

Molarity = Moles x 1000=>0.0225 moles x10000 = 22.5molesdm-3/M

1 cm3

 100%   =  22.5molesdm-3/M  

 68%   =   68 x 22.5 = 15.3M/ molesdm-3

  100

Method 2

Moles of HNO3 in 1000cm3 = Mass in 1000cm3 =>1.42 x1000

  Molar mass HNO3 63

  =22.5397 molesdm-3/M

 100%   =  22.5397 molesdm-3/M  

 68%   =   68 x 22.5397 = 15.327 molesdm-3

  100

(ii) the volume of ammonia gas at r.t.p used. (H=1.0, N=14.0, O=16.0, one mole of gas = 24 dm-3 at r.t.p)

Chemical equation

HNO3 (as)   + NH3 (g)  ->  NH4NO3 (as)

Mole ratio HNO3 (as): NH3 (g) = 1: 1

1 mole HNO3 (as) -> 24dm3 NH3 (g)

15.327 mole HNO3 (as) ->15.327 mole x 24 dm3 = 367.848dm3

1dm3

(iii) the number of crops which can be applied the fertilizer if each crop requires 4.0g.

HNO3 (aq)   + NH3 (g)   ->  NH4NO3 (aq)

 Molar mass NH4NO3 =80 g

 Mole ratio HNO3: NH4NO3 = 1: 1

Mass of HNO3 in 63.0 kg = 68% x 63 =42.84kg

 1 mole HNO3 (aq) =63g -> 80g NH4NO3


(42.84×1000) g
HNO3 (aq) -> (42.84×1000) g x 80

    63

  = 54400g

Mass of fertilizer =  54400g  = 13600 crops

Mass per crop   4.0

E. NITRATE (V) NO3 and NITRATE (III) NO2 Salts

Nitrate (V) /NO3 and Nitrate (III) /NO2 are salts derived from Nitric (V)/HNO3 and Nitric (III)/HNO2 acids. Both HNO3 and HNO2 are monobasic acids with only one ignitable hydrogen in a molecule.

Only KNO2, NaNO2 and NH4NO2 exist. All metallic nitrate (V) salts exist.

All Nitrate (V) /NO3 and Nitrate (III) /NO2 are soluble/dissolve in water.

(a)Effect of heat on Nitrate (V) /NO3 and Nitrate (III) /NO2 salts (Test for presence of Nitrate (V) /NO3 ions in solid state)

1. All Nitrate (III) /NO2 salts are not affected by gentle or strong heating except ammonium nitrate (III) NH4NO2.

Ammonium nitrate (III) NH4NO2 is a colourless solid that decompose to form Nitrogen gas and water.

Chemical equation

NH4NO2 (s)  -> H2O (l)  + N2 (g)


This reaction is used to prepare small amounts of Nitrogen in a school laboratory.

2. All Nitrate (V) /NO3 salts decompose on strong heating:

Experiment

Put ½ spatula full of sodium nitrate (V) into a test tube. Place moist blue/red litmus papers on the mouth of the test tube. Heat strongly when test tube is slanted.

Test the gases produced using glowing splint

Caution (i)
Wear safety gas mask and hand gloves

  (ii)Lead (II) nitrate (V) decomposes to Lead (II) oxide that reacts and fuses with the test tube permanently.


Repeat with potassium nitrate(V), copper(II) nitrate(V), Lead(II)nitrate(V), silver nitrate(V), Zinc nitrate(V), Magnesium nitrate(V) and Ammonium nitrate(V).

Observations

Cracking sound

Brown fumes/gas produced except in potassium nitrate (V) and Sodium nitrate (V)

Glowing splint relights/rekindles but feebly in Ammonium nitrate(V).

Black solid residue with copper(II) nitrate(V)

White residue/solid with sodium nitrate(V), potassium nitrate(V),silver nitrate(V), Magnesium nitrate(V)

Yellow residue/solid when hot but white on cooling with Zinc nitrate(V)

Brown residue/solid when hot but yellow on cooling with Lead(II)nitrate(V)

Explanation

1. Potassium nitrate(V) and Sodium nitrate(V) decomposes on strong heating to form potassium nitrate(III) and Sodium nitrate(III) producing Oxygen gas. Oxygen gas relights/rekindles a glowing splint.

Chemical equation.

2KNO3(s)  ->   2KNO2(s)   + O2 (g)

 2NaNO3(s)  ->   2NaNO2(s)   + O2 (g)

2.Heavy metal nitrate(V)salts decomposes to form the oxide, brown nitrogen (IV) oxide and Oxygen gas.

Copper(II)oxide is black.Zinc oxide is yellow when hot and white when cool/cold. Lead(II)oxide is yellow when cold/cool and brown when hot/heated.

Hydrated copper(II)nitrate is blue. On heating it melts and dissolves in its water of crystallization to form a green solution. When all the water of crystallization has evaporated,the nitrate(V)salt decomposes to black Copper(II)oxide and a mixture of brown nitrogen(IV)oxide gas and colourless Oxygen gas.

Chemical equation

2Cu(NO3)2(s) -> 2CuO (s) + 4NO2(g) + O2(s)  

2Ca(NO3)2(s) -> 2CaO (s) + 4NO2(g) + O2(s)

2Zn(NO3)2(s) -> 2ZnO (s) + 4NO2(g) + O2(s)

2Mg(NO3)2(s) -> 2MgO (s) + 4NO2(g) + O2(s)

2Pb(NO3)2(s) -> 2PbO (s) + 4NO2(g) + O2(s)

2Fe(NO3)2(s) -> 2FeO (s) + 4NO2(g) + O2(s)

 

Silver nitrate(V)and Mercury(II)nitrate decomposes to the corresponding metal and a mixture of brown nitrogen(IV)oxide gas and colourless Oxygen gas.

Chemical equation

2AgNO3 (s) -> 2Ag (s) + 2NO2(g) + O2(s)

Hg(NO3)2(s) -> Hg(l) + 2NO2(g) + O2(s)

The production/evolution of brown fumes of Nitrogen(IV)oxide gas on heating a salt is a confirmatory test for presence of NO3 ions of heavy metals

(b)Brown ring test (Test for presence of Nitrate(V) /NO3 ions in aqueous/ solution state)

Experiment

Place 5cm3 of Potassium nitrate(V)solution onto a clean test tube. Add 8 drops of freshly prepared Iron(II)sulphate(VI)solution. Swirl/ shake.

Using a test tube holder to firmly slant and hold the test tube, carefully add 5cm3 of Concentrated sulphuric (VI) acid down along the side of test tube.Do not shake the test tube contents.

Caution: Concentrated sulphuric (VI) acid is highly corrosive.

Observation.

Both Potassium nitrate(V)and freshly prepared Iron(II)sulphate (VI)do not form layers

On adding Concentrated Sulphuric(VI)acid,two layers are formed.

A brown ring is formed between the layers.

Image From EcoleBooks.com

Explanation

All nitrate(V)salts are soluble. They form a miscible mixture when added freshly prepared Iron(II)sulphate(VI)solution. Concentrated sulphuric(VI)acid is denser than the miscible mixture thus settle at the bottom.

At the junction of the layers, the acid reacts with nitrate(V)salts to form Nitric(V)acid/HNO3. Nitric(V)acid/HNO3 is reduced to Nitrogen (II)oxide by the Iron(II)sulphate(VI) salt to form the complex compound Nitroso-iron(II)sulphate(VI)/FeSO4.NO . Nitroso-iron(II)sulphate(VI) is brown in colour.It forms a thin layer at the junction between concentrated sulphuric (VI)acid and the miscible mixture of freshly prepared Iron(II) sulphate(VI) and the nitrate(V)salts as a brown ring.

Chemical equation

FeSO4(aq)   +  NO(g) -> FeSO4.NO(aq)

  (Nitroso-iron(II)sulphate(VI)complex)

The brown ring disappear if shaken because concentrated sulphuric (VI)acid mixes with the aqueous solution generating a lot of heat which decomposes Nitroso-iron(II)sulphate(VI)/FeSO4.NO to iron(II)sulphate(VI) and Nitrogen(II)oxide.

Chemical equation

FeSO4.NO(aq) ->FeSO4(aq)   +  NO(g) ->

Iron(II)sulphate(VI) solution is easily/readily oxidized to iron(III)sulphate(VI) on exposure to air/oxygen. The brown ring test thus require freshly prepared Iron(II) sulphate(VI) solution

(c)Devardas alloy test (Test for presence of Nitrate(V) /NO3 ions in aqueous/ solution state)

Experiment

Place 5cm3 of Potassium nitrate(V)solution onto a clean test tube. Add 5 drops of sodium hydroxide solution. Swirl/ shake. Add a piece of aluminium foil to the mixture.Heat.Test any gases produced using both blue and red litmus papers.

Observation. Inference

Effervescence/bubbles/fizzing

colourless gas that has a pungent smell of urine NO3

Blue limus paper remain blue

Red litmus paper turn red.

 

Explanation

The Devardas alloy test for NO3 ions in solution was developed by the Italian scientist Artulo Devarda(1859-1944)

When a NO3salt is added sodium hydroxide and aluminium foil, effervescence of ammonia gas is a confirmatory test for NO3ions.


 




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