HomeSS 2ChemistryCompounds of Nitrogen Chemistry SS 2 Second Term Lesson Notes Week 7

Compounds of Nitrogen Chemistry SS 2 Second Term Lesson Notes Week 7

Subject :Chemistry

Term : Second Term

Week :Week 7

Topic :

COMPOUNDS OF NITROGEN

Objectives:

  • Students will be able to identify the physical and chemical properties of nitrogen and its compounds.
  • Students will be able to describe the laboratory preparation of nitrogen (I) oxide and nitrogen (II) oxide.
  • Students will be able to explain the uses of nitrogen and its compounds in various industries and applications.

Materials:

  • Whiteboard and markers
  • Handouts on nitrogen and its compounds
  • Gas jars containing nitrogen and nitrogen (I) oxide
  • Bunsen burner and heat source
  • Reaction flask and catalyst for laboratory preparation of nitrogen (I) oxide and nitrogen (II) oxide
  • Safety goggles and gloves

WEEK SEVEN            

DATE: __________

TOPIC:COMPOUNDS OF NITROGEN

CONTENT

  • Oxides of Nitrogen
  • Ammonia: Preparation, Properties and Uses.
  • Trioxonitrate (V) acid: Preparation, Properties and Uses.

COMPOUNDS OF NITROGEN

Compounds of nitrogen are chemicals that contain the element nitrogen in their molecular structure. These compounds can be found in many everyday substances, including fertilizers, explosives, and pharmaceuticals. Here are some examples of common compounds of nitrogen:

  1. Ammonia (NH3): This is a colorless gas with a pungent odor. It is commonly used in cleaning products, refrigeration, and agriculture as a fertilizer. It is also used in the production of nylon and other synthetic fibers.
  2. Nitrogen dioxide (NO2): This is a reddish-brown gas with a sharp, pungent odor. It is a common air pollutant produced by the burning of fossil fuels, and exposure to high levels can be harmful to human health.
  3. Nitric oxide (NO): This is a colorless gas that is important in the body’s signaling processes. It is produced by cells in the blood vessels and helps to regulate blood pressure and blood flow.
  4. Nitrate (NO3): This is a compound that contains nitrogen and oxygen. It is commonly found in fertilizers and can also be present in drinking water. High levels of nitrate in drinking water can be harmful to infants and pregnant women.
  5. Nitroglycerin (C3H5N3O9): This is an explosive compound that is commonly used in construction and mining. It is also used in medicine to treat chest pain.
  6. Urea (CO(NH2)2): This is a compound that is commonly used as a fertilizer and in the production of plastics and resins. It is also a waste product produced by the body and is excreted in urine

 

Evaluation

  1. Which of the following is a compound of nitrogen? a) Sodium chloride b) Calcium carbonate c) Ammonia d) Hydrogen peroxide
  2. What is the chemical formula for nitrogen dioxide? a) NO2 b) N2O c) NH3 d) HNO3
  3. Which of the following is a common use of nitric oxide? a) Fertilizer b) Explosives c) Regulating blood pressure d) Cleaning products
  4. What can happen if there are high levels of nitrate in drinking water? a) Increased energy levels b) Improved immune system function c) Harmful effects on infants and pregnant women d) Reduced risk of cancer
  5. Which of the following compounds is commonly used to treat chest pain? a) Nitrogen dioxide b) Urea c) Nitric oxide d) Nitroglycerin
  6. What is the chemical formula for ammonia? a) NH4 b) H2SO4 c) NH3 d) CO2
  7. What is the common use of urea? a) Explosives b) Fertilizer c) Cleaning products d) Food additives
  8. Which of the following is a waste product produced by the body that contains nitrogen? a) Nitrate b) Nitric oxide c) Ammonia d) Nitroglycerin
  9. Which of the following is a common air pollutant that contains nitrogen? a) Nitrate b) Nitric oxide c) Urea d) Ammonia
  10. What is the chemical formula for nitroglycerin? a) C2H5OH b) C3H5N3O9 c) H2O2 d) NaCl

OXIDES OF NITROGEN

Oxides of nitrogen are a group of chemicals that contain nitrogen and oxygen. These compounds can be formed naturally, such as during lightning strikes and volcanic eruptions, or they can be produced by human activities, such as the burning of fossil fuels. Here are some examples of common oxides of nitrogen:

  1. Nitrogen monoxide (NO): This is a colorless gas that is produced naturally by lightning strikes and by certain bacteria in soil. It is also produced by the burning of fossil fuels in cars and power plants. In the atmosphere, it can react with other chemicals to form smog and acid rain.
  2. Nitrogen dioxide (NO2): This is a reddish-brown gas that is produced by the burning of fossil fuels. It is a common air pollutant that can cause respiratory problems and contribute to the formation of smog.
  3. Nitrous oxide (N2O): This is a colorless gas that is commonly known as “laughing gas” because of its use in dental anesthesia. It is also used as a propellant in whipped cream cans and as a greenhouse gas that contributes to climate change.
  4. Dinitrogen oxide (N2O2): This is a colorless gas that is used in rocket fuel and as an oxidizer in some chemical reactions.
  5. Dinitrogen trioxide (N2O3): This is a blue gas that is produced when nitrogen dioxide reacts with nitrogen monoxide. It is used in the production of nitric acid and other chemicals.
  6. Nitrogen pentoxide (N2O5): This is a white solid that is formed when nitrogen dioxide reacts with dinitrogen trioxide. It is used in the production of nitric acid and as a powerful oxidizing agent in chemical reactions.

Overall, oxides of nitrogen play important roles in natural and human processes, but their accumulation in the atmosphere can have negative impacts on human health and the environment

NITROGEN (I) OXIDE, N2O 

Nitrogen (I) oxide is known as laughing gas as it causes uncontrollable laughter when inhaled.

Nitrogen (I) oxide, also known as nitrous oxide or N2O, is a colorless and odorless gas that is used as an anesthetic in dentistry and surgery. As you mentioned, it is also known as “laughing gas” because inhaling it can cause feelings of euphoria and uncontrollable laughter.

Nitrous oxide is produced naturally in the atmosphere by bacteria and lightning, but it is also produced by human activities such as the burning of fossil fuels and the use of nitrogen-based fertilizers. In addition to its use as an anesthetic, it is also used as a propellant in whipped cream cans and in racing cars to increase engine power.

While nitrous oxide is relatively safe when used under medical supervision, prolonged exposure to high concentrations of the gas can lead to oxygen deprivation and other health issues. In addition, nitrous oxide is a potent greenhouse gas that contributes to climate change when released into the atmosphere. Therefore, it is important to use this gas in a responsible and safe manner

Evaluation

  1. What is the chemical formula for Nitrogen (I) oxide? a) NO b) N2O c) NO2 d) N2O2
  2. Which of the following is a common use of Nitrogen (I) oxide? a) As a fertilizer b) As a fuel c) As an anesthetic d) As a cleaning agent
  3. Why is Nitrogen (I) oxide also known as laughing gas? a) Because it causes a person to sneeze uncontrollably b) Because it causes a person to hiccup uncontrollably c) Because it causes a person to laugh uncontrollably d) Because it causes a person to cry uncontrollably
  4. How is Nitrogen (I) oxide produced naturally in the atmosphere? a) By the burning of fossil fuels b) By certain types of bacteria c) By volcanic eruptions d) By industrial activities
  5. What is the common use of Nitrogen (I) oxide as a propellant? a) In aerosol cans b) In cars c) In airplanes d) In boats
  6. What are the potential health risks of prolonged exposure to high concentrations of Nitrogen (I) oxide? a) Oxygen deprivation and other health issues b) Weight gain and other health issues c) Joint pain and other health issues d) Vision loss and other health issues
  7. Why is it important to use Nitrogen (I) oxide in a responsible and safe manner? a) Because it is a potent greenhouse gas that contributes to climate change b) Because it is highly explosive and can cause fires c) Because it can cause severe allergic reactions in some people d) Because it can cause skin irritation and other health issues
  8. What is the common use of Nitrogen (I) oxide in racing cars? a) To increase engine power b) To decrease engine power c) To improve fuel efficiency d) To reduce emissions
  9. Which of the following is a potential side effect of using Nitrogen (I) oxide as an anesthetic? a) Severe itching and rash b) Uncontrollable laughter and euphoria c) Difficulty breathing and chest pain d) Loss of taste and smell
  10. Which of the following is a recommended safety measure when using Nitrogen (I) oxide? a) Ensuring proper ventilation in the area b) Smoking cigarettes during use c) Storing it in a hot and humid area d) Mixing it with other chemicals to increase its effectiveness

LABORATORY PREPARATION 

Laboratory preparation refers to the process of preparing a chemical compound or solution in a laboratory setting. This can involve measuring and mixing specific quantities of chemicals, using appropriate equipment and safety precautions, and following established procedures. Here are some common methods of laboratory preparation:

  1. Solution preparation: This involves dissolving a solute in a solvent to create a solution. For example, a laboratory technician may prepare a 1M solution of sodium chloride by dissolving 58.44 grams of NaCl in 1 liter of water.
  2. Titration: This is a method used to determine the concentration of a substance in a solution. A standard solution of known concentration is added to the solution being tested until the reaction is complete, and the amount of standard solution required is measured.
  3. Distillation: This is a process used to separate two or more liquids with different boiling points. The mixture is heated to vaporize the component with the lower boiling point, which is then condensed and collected.
  4. Precipitation: This involves adding a reagent to a solution to cause a solid to form. The solid can then be filtered and dried to obtain a pure product. For example, a laboratory technician may prepare a sample of silver chloride by adding a solution of silver nitrate to a solution of sodium chloride.
  5. Extraction: This is a method used to separate a compound from a mixture by dissolving it in a solvent that is immiscible with the other components of the mixture. The solvent is then separated from the mixture and the compound is obtained by evaporating the solvent.

Overall, laboratory preparation involves careful planning and execution to ensure accurate and safe results. Proper equipment, safety precautions, and protocols should always be followed to minimize the risk of accidents and errors

LABORATORY PREPARATION OF NITROGEN (I) OXIDE, N2O

Nitrogen (I) oxide, also known as nitrous oxide or N2O, can be prepared in the laboratory by heating ammonium nitrate (NH4NO3) in a flask with a Bunsen burner or other heat source. Here are the steps involved in this process:

  1. Measure out 10 grams of ammonium nitrate using a balance and place it into a small round-bottomed flask.
  2. Attach a Bunsen burner to a retort stand and place the round-bottomed flask on a wire gauze on top of the Bunsen burner.
  3. Connect a gas collection apparatus to the flask. This can be a simple inverted funnel or a more complex gas collection tube.
  4. Heat the flask gently with the Bunsen burner until the ammonium nitrate starts to decompose. This reaction produces nitrogen (I) oxide gas and water vapor.
  5. Collect the nitrogen (I) oxide gas by displacement of water. This involves filling a gas collection apparatus with water and inverting it in a water bath, so that the open end of the apparatus is submerged in the water.
  6. As the nitrogen (I) oxide gas is produced, it will displace the water and rise up into the gas collection apparatus.
  7. Collect the nitrogen (I) oxide gas until the reaction is complete or until the desired amount of gas is obtained.
  8. Test the purity of the nitrogen (I) oxide gas using appropriate methods, such as gas chromatography or spectroscopy.

Laboratory preparation of nitrogen (I) oxide requires careful measurement of chemicals, appropriate equipment and safety precautions, and proper ventilation. It should only be done by experienced laboratory technicians who are familiar with the risks involved

The gas is prepared in the laboratory by thermal decomposition of ammonium trioxonitrate (V). Ammonium trioxonitrate (V) is not heated directly since the reaction is exothermic and may become uncontrollable leading to an explosion.  

KNO3(s) + NH4Cl(s) →   KCl(s) + NH4NO3(s)b NH4NO3(s) → 2H2O(g) + N2O(g)

PHYSICAL PROPERTIES 

  1. It is a colourless gas with a faint pleasant sickly smell and it has a sweetish taste. 
  2. It is fairly soluble in cold water. 
  3. It is 1.5 times denser than air.
  4. It is neutral to moist litmus paper. 
  5. Melting and boiling points: Nitrogen has a low melting point (-210°C) and boiling point (-196°C), which makes it a gas at room temperature and pressure.
  6. Density: Nitrogen is less dense than many common gases, such as oxygen and carbon dioxide. It has a density of 1.25 g/L at standard temperature and pressure (STP).
  7. State of matter: Nitrogen is a diatomic gas, which means it consists of two nitrogen atoms bound together in a molecule. At very low temperatures and high pressures, nitrogen can also exist as a liquid or solid.
  8. Non-toxicity: Nitrogen is non-toxic and does not react with most materials at normal temperatures and pressures. This makes it useful for storing and transporting delicate materials, such as food and pharmaceuticals.
  9. Inertness: Nitrogen is chemically inert, which means it does not readily react with other substances. This property makes it useful in many industrial applications, such as welding and laser cutting, where a non-reactive gas is needed to prevent oxidation and other unwanted reactions.
  10. Electrical conductivity: Nitrogen is a poor conductor of electricity, which makes it useful as an insulating gas in electrical equipment

 

 

CHEMICAL PROPERTIES 

  1. It decomposes on strong heating (about 600oC) to form nitrogen and oxygen. 2N2O(g) → O2(g)  +   2N2(g)
  1. It supports the combustion of any burning substance which is hot enough to decompose it. Mg(s)  +  N2O(g)  → MgO(s)  + N2(g)
  2. It is reduced to nitrogen by heated copper or iron Cu(s) + N2O(g)     → N2(g)  + CuO(s)
  3. Reactivity: Nitrogen is a relatively unreactive element and does not readily combine with other elements or compounds under normal conditions. However, it can react with highly reactive elements such as lithium and magnesium at high temperatures.
  4. Formation of nitrides: Nitrogen can react with some metals to form nitrides, such as lithium nitride (Li3N) and magnesium nitride (Mg3N2).
  5. Formation of oxides: Nitrogen can react with oxygen to form oxides, such as nitrogen monoxide (NO), nitrogen dioxide (NO2), and dinitrogen trioxide (N2O3). These oxides can contribute to air pollution and other environmental problems.
  6. Formation of ammonia: Nitrogen can react with hydrogen to form ammonia (NH3) through a process called Haber process. Ammonia is an important industrial chemical used as a fertilizer, refrigerant, and in the production of various other chemicals.
  7. Nitration reactions: Nitrogen is often used in organic chemistry to introduce nitro groups (-NO2) into organic molecules through a process called nitration. This process involves reacting nitrogen oxides with organic compounds under specific conditions

Evaluation

  1. What is the state of matter of nitrogen at room temperature and pressure? a) Solid b) Liquid c) Gas d) Plasma
  2. What is the density of nitrogen at standard temperature and pressure (STP)? a) 0.25 g/L b) 1.25 g/L c) 2.25 g/L d) 3.25 g/L
  3. What is the boiling point of nitrogen? a) -210°C b) -196°C c) -50°C d) 100°C
  4. What is the chemical formula for nitrogen monoxide? a) NO b) N2O c) NO2 d) N2O2
  5. What is the common use of nitrogen in the welding industry? a) To prevent oxidation b) To produce heat c) To increase pressure d) To generate electricity
  6. What is the process used to produce ammonia from nitrogen and hydrogen? a) Combustion b) Distillation c) Nitration d) Haber process
  7. What is the chemical formula for dinitrogen trioxide? a) NO b) N2O c) NO2 d) N2O3
  8. What is the chemical formula for magnesium nitride, a compound that can be formed by the reaction of nitrogen with magnesium? a) MgN2 b) Mg2N c) Mg3N2 d) Mg(NO3)2
  9. What is the property of nitrogen that makes it useful as an insulating gas in electrical equipment? a) High density b) Good conductivity c) Low reactivity d) High reactivity
  10. What is the process used to introduce nitro groups (-NO2) into organic molecules? a) Combustion b) Distillation c) Nitration d) Haber process

 

TEST FOR N2O

A glowing splinter is inserted into the gas jar containing the unknown gas. If the splinter is rekindled, the gas is either oxygen or nitrogen (I) oxide. If the gas has a pleasant smell and does not produce brown fumes with nitrogen (IV) oxide; then the gas is nitrogen (I) oxide.

This is a valid test to distinguish between oxygen and nitrogen (I) oxide, as both gases can rekindle a glowing splinter. However, it is important to note that this test alone cannot confirm the presence of nitrogen (I) oxide. The pleasant smell and lack of brown fumes with nitrogen (IV) oxide are additional characteristics that can suggest the presence of nitrogen (I) oxide, but further testing may be needed to confirm the identity of the gas. Other tests, such as the blue-black ring test or flame test, can be used to provide additional evidence for the presence of nitrogen (I) oxide. Additionally, it is important to use caution when handling nitrogen (I) oxide, as it is a powerful greenhouse gas and can be dangerous if not handled properly

There are several tests that can be used to identify the presence of nitrogen (I) oxide, or N2O. Here are two common tests:

  1. Blue-black ring test: This test involves adding a mixture of sulfuric acid and iron (II) sulfate to a sample suspected of containing N2O. The sample is then gently heated, causing any N2O present to decompose into nitrogen and oxygen. The nitrogen then reacts with the iron (II) sulfate to form a blue-black ring at the interface of the two layers.
  2. Flame test: This test involves heating a sample suspected of containing N2O in a flame. The heat causes the N2O to decompose into nitrogen and oxygen, which react with the flame to produce a characteristic yellow color.

Overall, these tests can be useful for identifying the presence of N2O, but they should be used with caution and only by experienced laboratory technicians who are familiar with the risks involved. N2O is a powerful greenhouse gas and can be dangerous if not handled properly. It should only be used in a laboratory setting with appropriate safety equipment and protocols in place

USES OF NITROGEN :

Nitrogen, which makes up 78% of Earth’s atmosphere, has many important uses in various industries and applications. Some common uses of nitrogen include:

  1. Fertilizer production: Nitrogen is an essential element for plant growth, and it is a key component of fertilizers used in agriculture.
  2. Food storage and packaging: Nitrogen gas can be used to create an inert atmosphere that inhibits the growth of bacteria and other microorganisms, extending the shelf life of packaged foods and preventing spoilage.
  3. Laser cutting: Nitrogen gas is used as a cutting and cooling agent in the laser cutting of metals, plastics, and other materials.
  4. Welding: Nitrogen gas can be used as a shielding gas during welding to prevent oxidation and other unwanted reactions.
  5. Semiconductor manufacturing: Nitrogen gas is used in the production of semiconductors and other electronic components to prevent contamination and improve yields.
  6. Cryogenics: Nitrogen is used as a coolant in cryogenic applications, such as in the storage and transportation of biological samples, medical specimens, and other materials at very low temperatures.
  7. Air conditioning and refrigeration: Nitrogen is used as a refrigerant in air conditioning and refrigeration systems, particularly in applications where the use of chlorofluorocarbons (CFCs) is prohibited.
  8. Fire suppression: Nitrogen gas can be used as a fire suppression agent in situations where the use of water or other extinguishing agents is not practical or effective.
  9. Oil and gas industry: Nitrogen is used in the oil and gas industry to increase pressure in wells and pipelines, to clean and maintain equipment, and to prevent the formation of explosive mixtures.
  10. Medical applications: Nitrogen is used in medical applications, such as in cryotherapy for the treatment of skin conditions, and in anesthesia as a diluent gas
  11. Nitrogen (I) oxide is used as anesthetic for minor surgical operations.

 

EVALUATION

  1. Describe the laboratory preparation of nitrogen (I) oxide.
  2. Describe a test to distinguish between nitrogen (I) oxide and oxygen gas.
  3. Which of the following tests can be used to distinguish between nitrogen and oxygen? a) Blue-black ring test b) Flame test c) Both A and B d) None of the above
  4. Which gas can cause uncontrollable laughter when inhaled? a) Nitrogen b) Oxygen c) Nitrogen (I) oxide d) Carbon dioxide
  5. What is the most common use of nitrogen in the food industry? a) As a flavoring agent b) As a refrigerant c) To create an inert atmosphere d) To prevent oxidation
  6. Which industry uses nitrogen as a coolant in cryogenic applications? a) Semiconductor manufacturing b) Oil and gas industry c) Food industry d) Medical industry
  7. What is the test used to confirm the presence of nitrogen (I) oxide? a) Blue-black ring test b) Flame test c) Pleasant smell and no brown fumes with nitrogen (IV) oxide d) Both A and C
  8. Which industry uses nitrogen to prevent contamination in the production of electronic components? a) Semiconductor manufacturing b) Oil and gas industry c) Food industry d) Medical industry
  9. Which gas is commonly used as a fire suppression agent? a) Oxygen b) Nitrogen c) Carbon dioxide d) Nitrogen (I) oxide
  10. Which of the following is not a use of nitrogen in the oil and gas industry? a) Increasing pressure in wells and pipelines b) Cleaning and maintaining equipment c) Preventing the formation of explosive mixtures d) Storing and transporting biological samples
  11. Which of the following is a property of nitrogen that makes it useful as a refrigerant? a) High density b) Good conductivity c) Non-toxicity d) Low reactivity
  12. What is the test used to distinguish between nitrogen and nitrogen (I) oxide? a) Blue-black ring test b) Flame test c) Pleasant smell and no brown fumes with nitrogen (IV) oxide d) None of the above

 

 

NITROGEN (II) OXIDE, NO  

Nitrogen (II) oxide, also known as nitrogen monoxide or NO, can be prepared in the laboratory by the reaction of nitric oxide (NO) and oxygen (O2) in the presence of a catalyst. Here are the steps involved in this process:

  1. Set up a reaction flask with a gas inlet and outlet, and attach it to a gas cylinder containing nitric oxide and a gas cylinder containing oxygen.
  2. Place a catalyst, such as copper or platinum, in the reaction flask.
  3. Heat the reaction flask with a Bunsen burner or other heat source to a temperature of around 600°C.
  4. Gradually introduce a mixture of nitric oxide and oxygen into the reaction flask through the gas inlet.
  5. The nitric oxide and oxygen will react with the catalyst to produce nitrogen monoxide and oxygen.
  6. Collect the nitrogen monoxide gas by displacement of water. This involves filling a gas collection apparatus with water and inverting it in a water bath, so that the open end of the apparatus is submerged in the water.
  7. As the nitrogen monoxide gas is produced, it will displace the water and rise up into the gas collection apparatus.
  8. Collect the nitrogen monoxide gas until the reaction is complete or until the desired amount of gas is obtained.
  9. Test the purity of the nitrogen monoxide gas using appropriate methods, such as gas chromatography or spectroscopy.

Overall, laboratory preparation of nitrogen (II) oxide requires careful measurement of chemicals, appropriate equipment and safety precautions, and proper ventilation. It should only be done by experienced laboratory technicians who are familiar with the risks involved. Nitrogen (II) oxide is a powerful greenhouse gas and can be dangerous if not handled properly

Nitrogen (II) oxide is prepared by reacting 50% trioxonitrate (IV) acid with copper. 

3Cu(s) + 8HNO3(aq)  →  3Cu(NO3)2(aq)  +  4H2O(l) + 2NO(g)

Some of the nitrogen (II) oxide gas reacts with oxygen in the flask to form brown fumes of nitrogen (IV) oxide which is dissolved in water as the gas is pass through water. 

LABORATORY PREPARATION

Nitrogen (II) oxide, also known as nitrogen monoxide or NO, can be prepared in the laboratory by the reaction of nitric oxide (NO) and oxygen (O2) in the presence of a catalyst. Here are the steps involved in this process:

  1. Set up a reaction flask with a gas inlet and outlet, and attach it to a gas cylinder containing nitric oxide and a gas cylinder containing oxygen.
  2. Place a catalyst, such as copper or platinum, in the reaction flask.
  3. Heat the reaction flask with a Bunsen burner or other heat source to a temperature of around 600°C.
  4. Gradually introduce a mixture of nitric oxide and oxygen into the reaction flask through the gas inlet.
  5. The nitric oxide and oxygen will react with the catalyst to produce nitrogen monoxide and oxygen.
  6. Collect the nitrogen monoxide gas by displacement of water. This involves filling a gas collection apparatus with water and inverting it in a water bath, so that the open end of the apparatus is submerged in the water.
  7. As the nitrogen monoxide gas is produced, it will displace the water and rise up into the gas collection apparatus.
  8. Collect the nitrogen monoxide gas until the reaction is complete or until the desired amount of gas is obtained.
  9. Test the purity of the nitrogen monoxide gas using appropriate methods, such as gas chromatography or spectroscopy.

Overall, laboratory preparation of nitrogen (II) oxide requires careful measurement of chemicals, appropriate equipment and safety precautions, and proper ventilation. It should only be done by experienced laboratory technicians who are familiar with the risks involved. Nitrogen (II) oxide is a powerful greenhouse gas and can be dangerous if not handled properly

PHYSICAL PROPERTIES 

  1. It is a colourless and poisonous gas. 
  2. It is a almost insoluble in water. 
  3. It is slightly denser than air. 
  4. It is neutral to litmus.
  5. Odor: Nitrogen (II) oxide has a slightly sweet odor, but at high concentrations it can cause irritation to the eyes and respiratory system. 
  6. Boiling point and melting point: Nitrogen (II) oxide has a boiling point of -152°C and a melting point of -163°C. 
  7. Solubility: Nitrogen (II) oxide is only slightly soluble in water, with a solubility of 0.027 g/100 mL at 25°C. 

Molecular weight: The molecular weight of nitrogen (II) oxide is 30.01 g/mol.

Overall, the physical properties of nitrogen (II) oxide make it a relatively stable gas at standard temperature and pressure, but also highlight the importance of handling the gas with care due to its toxicity and potential health hazards.

 

CHEMICAL PROPERTIES 

  1. It reacts readily with oxygen to form brown fumes of nitrogen (IV) oxide 2NO(g)   +    O2(g) → 2NO2(g)
  1. It decomposes on heating at high temperature to form equal volume of nitrogen and oxygen 2NO(s)→ N2(g) + O2(g)
  2. It is reduced to nitrogen by hot metals 2Cu(s)    +    2NO(g) →  2CuO(g)N2(g)
  3. It acts as reducing agent decolourizing acidified potassium tetraoxomanganate (VI) slowly 3MnO4-(aq)  +  4H+(aq)  + 5NO(g)   →  3Mn2+(aq)  +  5NO3-(aq) +  2H2O(l)
  4. Acid-base properties: Nitrogen (II) oxide is a neutral gas that does not react with most acids or bases. However, it can react with some strong acids to form nitrogen oxides and water.
  5. Formation of nitric oxide: Nitrogen (II) oxide can react with some metals, such as copper and silver, to form nitric oxide (NO). This reaction is often used to produce nitric oxide in the laboratory.
  6. Toxicity: Nitrogen (II) oxide is a poisonous gas that can cause harm to humans and animals if inhaled in large quantities. It is often used as a fumigant or anesthetic in medical settings, but care must be taken to avoid overexposure.
  7. Combustion: Nitrogen (II) oxide can support combustion in the presence of a fuel and oxygen. This can lead to the formation of nitrogen oxides and other harmful byproducts, making it important to handle the gas with care in industrial and laboratory settings

 

TEST FOR NO

  1. Using air: the gas jar containing the unknown gas is opened, if the gas turns reddish-brown, then the gas is NO.
  2. Using iron (II) tetraoxosulphate (VI): A solution of FeSO4 which has been acidified with a little dilute H2SO4 acid is poured into the gas jar containing the unknown gas. If the solution turns dark brown, then the gas is NO.

 

EVALUATION

  1. Give an equation to show the laboratory preparation of nitrogen (II) oxide.
  2. State TWO physical and TWO chemical properties of nitrogen (II) oxide.

 

NITROGEN (IV) OXIDE, NO2

LABORATORY PREPARATION

Nitrogen (IV) oxide is prepared by thermal decomposition of lead (II) trioxonitrate (V) because the nitrate does not contain water of crystallization which can interfere with the preparation. 

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

The gas mixture obtained is passed through a U- tube immersed in a freezing mixture. Nitrogen (IV) oxide liquefies as a green liquid (yellow if pure) in the tube while oxygen escapes out.

PHYSICAL PROPERTIES 

  1. It is a reddish – brown gas. 
  2. It has an irritating smell and is poisonous. 
  3. It turns damp blue litmus paper red and dissolves in water to form acidic solution.
  4. It liquefies into yellow liquid at 21oC. 
  5. It is much heavier than air.

 

CHEMICAL PROPERTIES 

  1. Nitrogen (IV) oxide exists mainly as dinitrogen (IV) oxide, N2O4 at low temperature. It decomposes on heating as follows. 

N2O4(g)       2NO2(g)     2NO(g) + O2(g)

Pale     Reddish       colourless 

yellow      brown

  1. It supports combustion as it decomposes on heating to nitrogen and oxygen 

2NO2(g) → N2(g)  + 2O2(g)

  1. It is reduced to nitrogen by reducing agent. 

4CU(s)  + 2NO(g) →  4CuO(s) + N2(g)

  1. It dissolves in water to form a mixture of dioxonitrate (III) and trioxonitrate (V) acids. It is a mixed acid anhydride. 

  H2O(l)  +  2NO2(g)   → HNO2(aq) + HNO3(aq)

  1.   It reacts with alkalis to form mixture of dioxonitrate (III) and trioxonitrate (V) salts 

2KOH(aq) + 2NO2(g) → KNO3(aq)   +   KNO2(aq)   +   H2O(l)

 

AMMONIA

Ammonia is a hydride of nitrogen. It is produced in nature when nitrogenous matter decays in the absence of air. Thus, traces of ammonia may be found in the atmosphere but being very soluble in water, it is dissolved by rain water and washed down into the soil.      

Ammonia (NH3) is a colorless gas with a pungent odor that is highly soluble in water. It consists of one nitrogen atom bonded to three hydrogen atoms and has a molecular weight of 17.03 g/mol. Ammonia is an essential building block in the production of fertilizers, as it provides a source of nitrogen for plants.

Ammonia is produced commercially by the Haber-Bosch process, which involves the reaction of nitrogen and hydrogen gases in the presence of a catalyst at high temperature and pressure. The resulting ammonia is then liquefied for transportation and storage.

In addition to its use in fertilizers, ammonia has many other applications. It is used in refrigeration systems as a refrigerant and as a cleaning agent in household and industrial settings. Ammonia is also a key component in the production of nylon, plastics, and explosives.

While ammonia has many useful properties, it can also be hazardous to human health if not handled properly. Exposure to high levels of ammonia can cause irritation and damage to the eyes, skin, and respiratory system. As such, proper precautions must be taken when working with or around ammonia

LABORATORY PREPARATION OF AMMONIA 

Ammonia is prepared in the laboratory by heating calcium hydroxide, Ca(OH)2 (slaked lime) with ammonium chloride. 

Ca(OH)2(s)+ 2NH4Cl(s) →  CaCl2(s) +2H2O(l)+2NH3(g). 

Ammonia is dried using calcium oxide, CaO. Ammonia being alkaline cannot be dried using Conc. H2SO4 or fused CaCl2, because they will react.

Ammonia can be prepared in the laboratory by the reaction of ammonium chloride (NH4Cl) and calcium hydroxide (Ca(OH)2), also known as the Haber process. Here are the steps for preparing ammonia in the laboratory:

Materials needed:

  • Ammonium chloride (NH4Cl)
  • Calcium hydroxide (Ca(OH)2)
  • Water
  • Heat source
  • Flask
  • Delivery tube
  • Collection bottle

Procedure:

  1. In a flask, mix equal amounts of ammonium chloride and calcium hydroxide. The reaction equation is:2NH4Cl + Ca(OH)2 → 2NH3 + CaCl2 + 2H2O
  2. Add enough water to the mixture to make a thick paste.
  3. Heat the flask gently using a heat source. The reaction will start and ammonia gas will be produced.
  4. Collect the ammonia gas by passing it through a delivery tube into a collection bottle. To make sure that no air is mixed with the ammonia, the delivery tube should be long enough to allow the gas to cool down before reaching the collection bottle.
  5. To test for the presence of ammonia, you can hold a piece of litmus paper near the opening of the delivery tube. Ammonia gas is basic and will turn red litmus paper blue.

Note: Ammonia gas is toxic and should be handled with care. Work in a well-ventilated area, wear gloves, goggles, and a mask to protect your eyes and respiratory system

INDUSTRIAL PREPARATION 

Ammonia is primarily produced industrially using the Haber-Bosch process, which involves the reaction of nitrogen and hydrogen gases in the presence of a catalyst at high temperature and pressure. The process consists of the following steps:

  1. The raw materials, nitrogen and hydrogen, are obtained from the atmosphere and from natural gas or other sources of hydrogen.
  2. Nitrogen gas is compressed and purified, and then mixed with hydrogen gas in a 3:1 ratio by volume.
  3. The mixture of nitrogen and hydrogen is heated to a temperature of about 450-500°C and passed over an iron catalyst.
  4. The nitrogen and hydrogen gases react to form ammonia gas according to the following equation:N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92.4 kJ/mol
  5. The reaction is exothermic, meaning that it releases heat. The heat generated is removed by cooling the reaction mixture and the ammonia gas that is produced is liquefied and stored for further use.
  6. The unreacted nitrogen and hydrogen gases are recycled back into the process, and the ammonia product is purified by distillation.

The Haber-Bosch process is a vital industrial process for the production of ammonia, which is used in the production of fertilizers, as well as in various other industries, including refrigeration, cleaning, and as a precursor to many important chemicals

Ammonia is manufactured from nitrogen and hydrogen by the Haber process. It involves mixing nitrogen and hydrogen in ratio 1:3 by volume. The reaction is reversible so special conditions listed below are required for optional yield of ammonia. 

  1. Finely divided iron catalyst is used 
  2. Temperature of about 450Oc is used 
  3. Pressure of about 200atm is used. 

The yield is about 15% under this condition 

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

 

PHYSICAL PROPERTIES 

  1. It is a colorless gas with a characteristic choking smell. 
  2. Ammonia in large quantity is poisonous as it affects respiratory muscles. 
  3. It is the only known alkaline gas. 
  4. It is about 1.7 times less dense than air. 
  5. Solid ammonia melts at -34.4OC and the liquid boils at -77.7OC. 

 

CHEMICAL PROPERTIES 

  1. Ammonia burns readily in oxygen to form water vapor and nitrogen 

4NH3(g)   +  3O2(g)   → 6H2O(g)  +  2N2(g)

  1. Ammonia reacts as reducing agents reacting with 
  2. Copper II oxide 

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

  1. Chlorine 

3Cl2(g) +  8NH3(g) →      6NH4Cl(s) +  N2(g)

  1. Ammonia reacts with carbon IV oxide to form Urea and water vapour. 

2NH3(g)   +   CO2(g) →  (NH2)2 CO(s) + H2O(l)

urea 

  1. Ammonia reacts with acid to form ammoniums salts.   

2NH3(g)  +  H2SO4(g)  →  (NH4)2SO4(s)

Ammonia (NH3) is a colorless gas with a pungent odor that is highly soluble in water. Here are some of the important physical properties of ammonia:

  1. Boiling point: The boiling point of ammonia is -33.34°C, which is significantly lower than the boiling points of other gases in the air, such as oxygen and nitrogen.
  2. Melting point: The melting point of ammonia is -77.73°C, which is lower than the melting point of water.
  3. Density: The density of ammonia gas is lower than that of air. At standard temperature and pressure (STP), the density of ammonia gas is 0.73 g/L.
  4. Solubility: Ammonia is highly soluble in water, with a solubility of about 700 g/L at STP. This makes it easy to dissolve and transport in water, which is why it is often found in natural waters and used in fertilizer applications.
  5. Odor: Ammonia has a pungent odor that is easily detectable by most people, even at low concentrations.
  6. Flammability: Ammonia is not flammable but it can support combustion of other materials.
  7. Reactivity: Ammonia is a highly reactive gas and can react with a wide range of other chemicals, including acids, halogens, and metals.
  8. Toxicity: Ammonia gas is toxic to humans and animals if inhaled in high concentrations. It can cause severe irritation and damage to the respiratory system and eyes

 

TEST FOR AMMONIA

Ammonia has a choking smell. It can be confirmed using:

  1. Litmus paper: Damped red litmus is dipped into the gas jar containing the unknown gas. If the litmus paper turns blue, then the gas is ammonia.
  2. Hydrochloric acid: a glass rod is dipped in concentrated HCl and then inserted in the gas jar containing the unknown gas. If white fumes are observed on the glass rod, then the gas is ammonia.

 

USES OF AMMONIA 

  1. Ammonia is used in the manufacture of trioxonitrate (V) acid and Sodium trioxocarbonate (IV) by the Solvay process.
  2. Liquid ammonia is used as a refrigerant.
  3. Aqueous ammonia is used in softening temporary hard water.
  4. Aqueous ammonia is also used in laundries as a solvent for removing grease and oil stains.

 

EVALUATION 

  1. Briefly describe the laboratory preparation of ammonia. 
  1. StateTWO physical and THREE chemical properties each of ammonia.

 

TRIOXONITRATE (V) ACID, HNO3

LABORATORY PREPARATION 

Trioxonitrate (V) acid is a volatile acid and it is prepared in the laboratory by its displacement from any trioxonitrate salt by concentrated H2SO4 which is less volatile. Trioxonitrate (V) of potassium or sodium is usually used because they are cheap. 

KNO3(s)    +   H2SO4(aq) →  KHSO4(aq)  +  HNO3(aq)

 

NOTE: An all glass apparatus must be used in this preparation because the hydrogen trioxonitrate (V) acid vapour will attack cork or rubber.

INDUSTRIAL PREPARATION 

Trioxonitrate (V) acid is obtained by the catalytic oxidation of ammonia: 

–    Ammonia is treated with excess air using Platinum-rhodium catalyst at 700oC to produce nitrogen (II) oxide (96% conversion of NH3 is obtained) 

4NH3(g)    +    5O2(g) → 4NO(g)  +  6H2O(g)

–    Nitrogen (II) oxide formed is cooled and mixed with excess air to produce nitrogen (IV) oxide.

2NO(g)    +    O2(g) →  2NO2(g) 

–    Nitrogen (IV) oxide formed is dissolved with excess air in hot water to yield trioxonitrate (V) acid solution of up to 50% concentration.  

4NO2(g)  +  2H2O(l)  +    O2(g) →  4HNO3(aq)

 

PHYSICAL PROPERTIES 

  1. The pure acid is a colourless fuming liquid with sharp choking smell. The acid turns yellow due to its decomposition to nitrogen (IV) oxide which redissolves in the acid. 
  2. The pure acid boils at 860C and melts at -47oC 
  3. The density of the pure acid is 1.52 gcm-3
  4. The pure acid is miscible with water in all properties and forms constant boiling mixture with it at 121oC 
  5. The concentrated form of the acid is corrosive. 
  6. The dilute acid turns blue litmus red.

 

CHEMICAL PROPERTIES 

  1. As an acid it neutralizes bases and alkalis to form metallic trioxonitrate (V) and water only 

NaOH(aq) + HNO3(aq) →  NaNO3(aq) + H2O(l)

  1. As an acids it reacts with metallic trioxocarbonate (IV) to liberate 

carbon (II) oxide  

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

  1. Unlike other acids, it rarely gives out hydrogen with metals except when very dilute solution is reacted with Ca, Mg or Mn. 
  2. As an oxidizing agent, it reacts with non – metal to form the corresponding oxides of the non – metals. 

S(s) + 6HNO3(aq)   →  H2SO4(aq) + 2H2O(l) +    6NO2(g)

  1. As an oxidizing agent, it oxidizes Cu, Pb, Hg and Ag to yield the

respective trioxonitrate (V) and nitrogen (IV) oxide if concentrated, but nitrogen (II) oxide if the concentration is moderate. 

Aluminum and iron are not oxidized to their oxides by concentrated HNO3(aq) due to formation of a surface coating of oxide which is passive do not allow further reaction with the metals. Aluminum or iron lined container can be used to transport concentrated HNO3(aq)

  1. As an oxidizing agent, it oxidizes hydrogen sulphide to sulphur 

H2S(g) + 2HNO3(aq)  → S(s)  + 2H2O(l)  +  2NO2(g)

  1. As an oxidizing agent, it oxidizes iron (II) salts to iron (III) salts 

6Fe2+(aq) + 8H+(aq) + 2NO3–(aq)      →   6Fe3+(aq) + 4H2O(l) + 2NO(g)

 

USES 

  1. It is used as an acid, oxidizing agent and nitrating agent in the laboratory. 
  2. It is used in nylon production and Terylene.
  3. It is used as rocket fuel.
  4. It is used in production of fertilizers, dyes, drugs and explosives.

 

GENERAL EVALUATION/REVISION

  1. Describe the laboratory preparation of trioxonitrate (V) acid.
  2. Write TWO equations of reactions in which trioxonitrate (V) is acting as an acid. 
  3. Write an equation to show the reaction of nitrogen (IV) oxide as a mixed anhydride.
  4. Describe the electrolysis of CuSO4 solution using platinum electrodes.
  5. Classify the following oxides: CuO, Na2O, PbO, NO2, N2O

 

READING ASSIGNMENT

New School Chemistry for Senior Secondary Schools by O. Y. Ababio (6th edition), pages 406-409, 411-419. 

 

WEEKEND ASSIGNMENT

SECTION A: Write the correct potion ONLY.

  1. Pure trioxonitrate (V) acid is colourless but the product of its laboratory preparation is yellow because of the presence of dissolved a. N2O b. NO c. NO2 d. NH3
  2. Common laboratory drying agents are not used for drying ammonia because a. ammonia is alkaline b. ammonia forms complexes with them c. ammonia reacts with them and disappears into products d. ammonia is highly soluble in water 
  3. Ammonia has relatively high boiling point when compared with other similar compounds because a. ammonia is stable b. ammonia is easily liquefied c. ammonia has hydrogen bonding d. ammonia is soluble in water.
  4. Aqueous ammonia solution used in the laboratory is referred to as aqueous ammonia and not ammonium hydroxide because a. ammonia dissolves in water without forming bond b. ammonia solution easily decomposes and liberated free ammonia when the temperature of the room rises leaving water in the bottle c. bond between ammonia and OH– of water is weakly acidic d. ammonia is less dense than air.
  5. Which of the following metals would be in passive state when treated with concentrated HNO3? a. Zinc b. Sodium c. Tin d. Iron  

 

SECTION B

  1. Give reason for the following
  1. The flask used for the laboratory preparation of ammonia is mounted in a slanting position
  2. An all glass apparatus is used for the laboratory preparation of trioxonitrate (V) acid. 
  1. Give an example of a reaction in which ammonia behaves as a
  2. reducing agent b. base c. precipitating agent

Ammonia (NH3) is a highly reactive compound with a variety of chemical properties. Here are some of the important chemical properties of ammonia:

  1. Basicity: Ammonia is a weak base and can react with acids to form salts. It reacts with water to produce ammonium ions (NH4+) and hydroxide ions (OH-).NH3 + H2O ⇌ NH4+ + OH-
  2. Acidity: Ammonia can also act as a weak acid in the presence of strong bases. It can react with strong bases such as sodium hydroxide (NaOH) to form sodium amide (NaNH2) and water.NH3 + NaOH → NaNH2 + H2O
  3. Reducing agent: Ammonia is a powerful reducing agent and can donate electrons to other molecules. For example, it can reduce silver ions (Ag+) to metallic silver (Ag).2Ag+ + 2NH3 → 2Ag + 2NH4+
  4. Complex formation: Ammonia can form complexes with metal ions, such as copper (Cu2+), iron (Fe2+), and nickel (Ni2+). These complexes have important industrial and biological applications.Cu2+ + 4NH3 → [Cu(NH3)4]2+
  5. Combustibility: Ammonia is not combustible but can support the combustion of other materials. In the presence of oxygen, ammonia can react to form nitrogen oxides (NOx), which are a major source of air pollution.4NH3 + 5O2 → 4NO + 6H2O
  6. Toxicity: Ammonia gas is toxic to humans and animals if inhaled in high concentrations. It can cause severe irritation and damage to the respiratory system and eyes. Ammonia can also be harmful to aquatic life if released into waterways

 

Test For Ammonia

There are several tests that can be performed to detect the presence of ammonia (NH3). Here are two common tests:

  1. The damp red litmus paper test: Ammonia is a weak base, and when it is introduced to a damp piece of red litmus paper, it will turn the paper blue. This is because the ammonia reacts with the water on the paper to form ammonium hydroxide, which is an alkaline solution.
  2. The white fumes test: When a small amount of concentrated hydrochloric acid (HCl) is added to a solid or liquid sample suspected of containing ammonia, white fumes will be produced if ammonia is present. The white fumes are actually ammonium chloride (NH4Cl), which is formed by the reaction of ammonia with the hydrochloric acid.NH3 + HCl → NH4Cl

These tests are relatively simple and can be performed quickly in a laboratory setting. However, it is important to note that ammonia is toxic and should be handled with care. Proper ventilation and personal protective equipment should be used when working with ammonia

 

 

Uses of Ammonia

Ammonia (NH3) has numerous uses in various industries. Here are some of the most common uses of ammonia:

  1. Fertilizer production: Ammonia is used as a key building block in the production of nitrogen-based fertilizers, which provide a source of nitrogen for plants.
  2. Refrigeration: Ammonia is widely used as a refrigerant in industrial refrigeration systems, such as those used in food processing and cold storage.
  3. Cleaning: Ammonia is used as a cleaning agent in household and industrial settings due to its ability to dissolve grease and oils.
  4. Textile production: Ammonia is used in the production of nylon, which is a synthetic fiber used in clothing and other textiles.
  5. Water treatment: Ammonia is used in the treatment of municipal water supplies and wastewater to control the growth of harmful bacteria.
  6. Chemical production: Ammonia is used as a precursor to many important chemicals, including ammonium nitrate, which is used as a fertilizer and explosive.
  7. Pharmaceuticals: Ammonia is used in the production of a variety of pharmaceuticals, including antibiotics, sedatives, and analgesics.
  8. Metal treatment: Ammonia is used in the treatment of metal surfaces to remove oxide layers and promote adhesion of coatings.

Overall, ammonia is a versatile compound with a wide range of uses in various industries

 

Evaluation

  1. What is the chemical formula for ammonia? a) NH b) NH2 c) NH3 d) N2H4
  2. Which of the following is NOT a physical property of ammonia? a) Boiling point of -33.34°C b) Highly soluble in water c) Pungent odor d) Strong reducing agent
  3. What is the primary use of ammonia in the agricultural industry? a) As a refrigerant b) In the production of nylon c) In the production of pharmaceuticals d) As a key building block in nitrogen-based fertilizers
  4. What is the name of the industrial process used to produce ammonia? a) Haber-Bosch process b) Fisher-Tropsch process c) Solvay process d) Contact process
  5. Which of the following is a test used to detect the presence of ammonia? a) Blue litmus paper test b) Red litmus paper test c) White fumes test d) Benedict’s test
  6. What happens when ammonia is introduced to a damp piece of red litmus paper? a) The paper turns red b) The paper turns blue c) The paper does not change color d) The paper dissolves
  7. What is the boiling point of ammonia? a) -273.15°C b) -196°C c) -78°C d) -33.34°C
  8. Which of the following is a common industrial use of ammonia? a) Production of toothpaste b) Production of glass c) Production of paper d) Refrigeration
  9. Which of the following is a chemical property of ammonia? a) High solubility in water b) Pungent odor c) Reacts with acids to form salts d) Boiling point of -33.34°C
  10. What is the primary danger associated with exposure to high levels of ammonia? a) Fire hazard b) Explosive hazard c) Skin irritation d) Respiratory damage
  11. What is ammonia, and what is its chemical formula?
  12. What are the physical properties of ammonia?
  13. What is the primary industrial use of ammonia?
  14. How is ammonia produced industrially?
  15. How can the presence of ammonia be detected in a laboratory setting?
  16. What is the damp red litmus paper test used for?
  17. What is the boiling point of ammonia?
  18. What are some of the chemical properties of ammonia?
  19. What are some of the safety concerns associated with handling ammonia?
  20. What are some of the environmental impacts of ammonia use in industry and agriculture?

Lesson Plan Presentation: Nitrogen and Its Compounds

 

Grade Level: SS2

 

Introduction (10 minutes):

  • Begin by asking students what they know about nitrogen and its compounds.
  • Use the whiteboard to create a mind map or concept web of student responses.
  • Introduce the objectives for the lesson and explain the importance of understanding nitrogen and its compounds in various fields.

Main Lesson (40 minutes):

  • Discuss the physical and chemical properties of nitrogen, including its density, solubility, and neutrality to litmus paper.
  • Demonstrate the reactivity of nitrogen (I) oxide by using a glowing splinter test or blue-black ring test to confirm its presence.
  • Discuss the laboratory preparation of nitrogen (I) oxide and nitrogen (II) oxide, highlighting the importance of proper handling and safety precautions.
  • Describe the uses of nitrogen and its compounds in various industries, such as fertilizer production, food storage and packaging, welding, and semiconductor manufacturing.
  • Show examples and visuals of nitrogen and its compounds in use in each industry.

Activity (20 minutes):

  • Divide students into small groups and provide each group with a gas jar containing either nitrogen or nitrogen (I) oxide.
  • Ask students to perform a test to distinguish between the two gases, such as a glowing splinter test or blue-black ring test.
  • Encourage students to record their observations and explain their reasoning for identifying the gas in their jar.

Conclusion (10 minutes):

  • Have students share their observations and conclusions from the activity.
  • Recap the main points of the lesson and review the physical and chemical properties of nitrogen and its compounds, as well as their uses in various industries and applications.
  • Allow time for any remaining questions or discussion on the topic.

Assessment:

  • Monitor student participation in the group activity and assess their understanding of the physical and chemical properties of nitrogen and its compounds.
  • Administer a quiz or written assignment to test students’ comprehension of the material covered in the lesson.

Weekly Assessment /Test

  1. What is the most abundant gas in Earth’s atmosphere? a) Oxygen b) Nitrogen c) Carbon dioxide d) Methane
  2. What is the laboratory preparation method for nitrogen (I) oxide? a) Reaction of nitrogen and oxygen at high temperature b) Reaction of nitric oxide and oxygen in the presence of a catalyst c) Reaction of nitrogen and hydrogen d) Reaction of nitrogen (II) oxide and water
  3. What is the test used to distinguish between nitrogen and nitrogen (I) oxide? a) Glowing splinter test b) Blue-black ring test c) Flame test d) Acid-base titration
  4. What is the chemical formula for nitrogen (I) oxide? a) NO b) NO2 c) N2O d) N2O2
  5. What is the chemical formula for nitrogen (II) oxide? a) NO b) NO2 c) N2O d) N2O2
  6. What is the common use of nitrogen in the food industry? a) As a flavoring agent b) As a refrigerant c) To create an inert atmosphere d) To prevent oxidation
  7. What is the physical state of nitrogen (II) oxide at room temperature? a) Solid b) Liquid c) Gas d) Plasma
  8. What is the chemical equation for the decomposition of nitrogen (I) oxide at high temperature? a) 2NO(g) + O2(g) → 2NO2(g) b) 2NO(s) → N2(g) + O2(g) c) Cu(s) + 2NO(g) → CuO(g) + N2(g) d) 3MnO4-(aq) + 4H+(aq) + 5NO(g) → 3Mn2+(aq) + 5NO3-(aq) + 2H2O(l)
  9. Which of the following is a use of nitrogen (II) oxide in the medical industry? a) As a coolant in cryogenic applications b) As a fumigant c) As a laser cutting agent d) As a fertilizer
  10. What is the chemical formula for nitrogen gas? a) N b) N2 c) NO d) NO2

Recommend Posts :

  1. The Periodic Trends Chemistry SS 2 Second Term Lesson Notes Week 2
  2. SS 2 Chemistry Third Term Examination
  3. THE PERIODIC TABLE
  4. IONIC THEORY
  5. CHARACTERISTICS PROPERTIES OF TRANSITION ELEMENTS
  6. CHEMICAL REACTIONS BASIC CONCEPTS: REACTANTS AND PRODUCTS
  7. Exothermic and Endothermic Reactions
  8. CHEMICAL EQUILIBRIUM
  9. Oxygen and Its Compounds
  10. CHLORINE AND OTHER HALOGENS
Tags: