Chemistry

Key unit competency: Compare and contrast the properties of Group 15 elements and their compounds, in relation to their position in the Periodic Table.

Learning objectives

By the end of this unit, the learner will be able to:

•Describe the physical properties of Group 15 elements.

•Describe the variation in the metallic and non-metallic character of Group 15 elements.

•Explain recall the physical properties of the allotropes of phosphorus.

•Describe the chemical reactions of nitrogen compounds.

•Describe the impact of nitrogen oxides to the environment.

•Describe the industrial preparation of ammonia and nitric acid.

•Explain the reactions of nitric acid with metals and non-metals.

•Describe the chemical properties of phosphorus compounds.

•State the uses of the group 15 elements and its compounds

10.1. Physical properties of group 15 elements and the relative inertness of nitrogen

Activity 10.1

In pairs:

1. Assign the physical state for each of the elements in group 15

2. Explain what is meant by the term “metallic character”.

3. Classify each element in this group as metal, non-metal or metalloid

.4. Study the following figure carefully and answer the questions that follow

a. Identify the molecule represented in the figure.

b. What type of bond is there in the molecule?

c. Suggest if the bond is strong or weak.

10.1.1.Physical properties of group 15 elements

a. Physical stateThis group consists of nitrogen (7N), phosphorus (15P), arsenic (33As), antimony (51Sb), and bismuth (83Bi). This group consists of atoms having 5 electrons in their outer energy level. Two of the electrons are in the s sub-shell, with 3 unpaired electrons in the p sub-shell: ns2np3 (px1py1pz1). Group 15 Elements often form covalent compounds, usually with the oxidation numbers +3 or +5.

All the elements of this group are polyatomic. Dinitrogen is a diatomic gas while all others are solids.

•Nitrogen is a colorless gas and is the major component of air (78%).

•Phosphorus has allotropic forms and is avery reactive non-metallic element,

•Arsenic, antimony (their oxides are amphoteric) and bismuth (its oxide is basic) are solidin nature.

Table 10.1: Physical properties of group 15 elements

b. Metallic character

Down group 15 elements, the atomic radius increases which makes the outermost electron to be less attracted by the nucleus as you move down the group. Therefore, less energy is required to remove the outermost electron, which results in the increase in the metallic character down the group. This results also in decreasing of ionization energy down the group.

Nitrogen and phosphorous are non-metals, with the metallic properties first appearing in arsenic and increasing down the group. Arsenic and antimony are metalloids. Bismuth is a metal.

10.1.2. The relative inertness of nitrogen

The most striking aspect of nitrogen chemistry is the inertness of N2 itself. Nearly four-fifth of the atmosphere consists of nitrogen, and the other fifth is nearly all oxygen, a very strong oxidizing agent. Nevertheless, the searing temperature of a lightning bolt or very high temperature in a vehicle motor combustion is required for significant amounts of atmospheric nitrogen oxides to form. Thus, even though N2 is inert at moderate temperatures, it reacts at high temperatures with H2, Li, Group 2 members, B, Al, C, Si, Ge, O2, and many transition elements. In fact, nearly every element in the periodic table forms bonds to N, but at high temperature.

Nitrogen is found as a diatomic molecule (N2) in nature. There is a triple bond between the two nitrogen atoms. This explains the very high strength of the nitrogen-nitrogen triple bond (N≡N) that requires high energy, 942 kJmol-1, to be broken. For comparison, the single bond (N-N) and the double bond (O=O) require 247kJmol- and 498 kJmol- respectively.

Checking Up 10.1

1. Briefly describe how each of the following factors varies in group 15 elements:

a) Atomic radius

b) Electron affinity

c) Melting point

d) First ionization energy

2. Explain the following observations:

a) In group 15 of the periodic table, metallic character increases as you move down the group.

b) The atomic radii of two elements A and B from group 15 are 0.121nm and 0.141nm respectively. Identify the element with more metallic character. Justify your answer.

10.2. Reactions of group 15 elements

Activity 10.2

1. Explain what is meant by the term:

a. Hybridisation

b. Electronic configuration

c. Briefly explain what is meant by:

d. Acidic oxide

e. Basic oxide

f. Amphoteric oxide

2. Respond by True or False and justify

a. Nitrogen is the only member of the group that readily forms multiple covalent bonds.

b. Nitrogen is the most electronegative member of group 15 elements and the only one that can be involved in the formation of hydrogen bond

c. Nitrogen is the only group member that has an oxidation state -3.

d. Only nitrogen form oxides of the form: E2O3 and E2O5

All group 15 elements exhibit a common valency of three. They can complete their octet structure in chemical combination by gaining three electrons.

However, with the exception of nitrogen, group 15 elements have vacant d-orbitals which they use to expand their octet to form compounds with a valency of five. For instance phosphorous has a covalency of 5 due to availability of easily accessible empty d orbitals which can be used for sp3d hybridization that allows it to have 5 unpaired electrons. Consider phosphorous, atomic number 15.

Ground state electronic configuration of P: 1s2 2s2 2p6 3s2 3p3 3d0 (three unpaired electrons)

After sp3d hybridization, the 5 electrons previously in 3s23p3 are redistributed in the five hybrid orbitals sp3d:

Hence phosphorus can now accommodate 5 more electrons or create 5 covalent bonds.

The same applies for any element in the group, such as Bi, that participates in a five covalency bonding.

Reaction with oxygen

All group 15 elements form two types of oxides: E2O3 and E2O5. The oxide in the higher oxidation state of the element is more acidic than that of lower oxidation state.Their acidic character decreases down the group. Group 15 oxides of the form E2O3 of nitrogen and phosphorus are purely acidic, those of arsenic and antimony amphoteric and that of bismuth predominantly basic.

Nitrogen gas does not react with air under normal conditions. Nitrogen (II) oxide is produced when nitrogen combines with oxygen. This is further oxidized into nitrogen (IV) oxide:

Reaction with water

Elements in this group do not tend to react with water. Nitrogen gas does not react with water. Phosphorous does not react with cold water. However, phosphorous vapour reacts with steam at high temperatures.

Reaction with chlorine

White phosphorous readily ignites in chlorine to form the trichloride and in excess of chlorine, the pentachloride is produced. Red phosphorous behaves similarly when heated:

Reaction with hydrogen

All Group 15 elements form hydrides of the type EH3, where E = N, P, As, Sb or Bi. The hydrides show regular gradation in their properties. The stability of hydrides decreases from NH3 to BiH3 which can be observed from their bond dissociation energy. Consequently, the reducing character of the hydrides increases down the group. NH3 is only a mild reducing agent while BiH3 is the strongest reducing agent amongst all the hydrides of the Group 15. Basicity also decreases in the order:

Under normal conditions the nitrogen gas is very inert, but, Nitrogen reacts with Hydrogen at a high temperature (about 500oC) and extremely high pressure (200-1000atm) over a catalyst (consisting of finery divided iron (Fe) and aluminum oxides a promoter). These are conditions used in the manufacture of ammonia, NH3:

Reaction with metals

All group 15 elements react with metals to form their binary compounds exhibiting –3 oxidation state, such as, Ca3N2(calcium nitride), Ca3P2(calcium phosphide), Na3As2(sodium arsenide), Zn3Sb2(zinc antimonide) and Mg3Bi2(magnesium bismuthide).

Nitrogen combines with many metals to form metal nitride:

Checking Up 10.2

1. Group 15 elements form compounds by gaining three electrons, for example in the formation of ionic nitrides such as Li3N, Ca3N2, Mg3N2 and AlN, in which nitrogen forms N3-. Other large atoms do not form M3-. Explain why.

2. State how the stability of +5 oxidation state varies in group 15.

3. The atomic number of an element A is 33.

a. Write the electronic configuration of A using s,p,d, notation.

b. Write the molecular formula of all possible oxide of A.

c. i) State whether each oxide of A you have given in (b) is acidic, basic, neu-tral, or amphoteric and justify.

ii) Write the equation of reaction to illustrate your answer.

10.3. Ammonia and nitric acid

Activity 10.3 (a)

Experiment: Laboratory preparation of ammonia

Materials and chemicals

Round bottom flaskor hard glass test tube, U-tube, 3 corks, 10 grams of calcium hydroxide, gas jars, bent delivery tube and straight delivery tube, 5 grams of ammonium chloride on a watch glass and calcium oxide lumps.

Procedure

1. Set up the apparatus as shown in the diagram, with the chemicals indicated. Do not start heating yet.

2. When everything is in position, heat the hard flaskandcollect several gas jarsof ammonia. Cover each jar with a glass slip and keep the jars for other experiments.

Activity evaluation questions

1. Record your observations

2. Write a balanced equation of the reaction that take place.

10.3.1. Laboratory preparation of ammonia and nitric acid

a. Laboratory preparation of ammonia

Ammonia is a covalent compound, consisting of nitrogen bonded to three hydrogen atoms. It exists as a colourless gas at room temperature and it is naturally produced during the decaying of nitrogenous organic compounds such as proteins. Ammonia has a characteristic pungent odour.It is less dense than air and thus collected by upward delivery method. In the laboratory it is prepared by heating a mixture of any ammonium salt and an alkali.

Examples

10.3.2. Industrial production of ammonia and nitric acid

Production of ammonia (Haber process)

The Haber process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic.

This reaction, which is reversible, requires special conditions for an optimum yield of ammonia

•Pressure : 200 –1000 atmospheres

•Temperature: 450- 500 ̊C

•Catalyst: finely divided iron (Fe) catalyst, aluminium oxide which is added to the catalyst to improve its performance (a promoter). It makes it more porous and this provides a high surface area to the reaction and potassium oxide

The ammonia formed can either be liquefied or dissolved in water to separate it from the mixture.

i. Uses of ammonia

Agricultural industries are the major users of ammonia. Ammonia and urea are used as fertilizer, as very valuable source of nitrogen that is essential for plant growth.

Ammonia and urea are used as a source of protein in livestock feeds for ruminating animals such as cattle, sheep and goats.

Ammonia can also be used as a pre-harvest cotton defoliant, an anti-fungal agent on certain fruits and as preservative for the storage of high-moisture corn.

The pulp and paper industry uses ammonia for pulping wood and as casein dispersant in the coating of paper.

The food and beverage industry uses ammonia as a source of nitrogen needed for yeast and microorganisms involved in the fermentation process.

ii. Environmental impact for industrial production of ammonia

Making ammonia using the Haber process requires a lot of energy, which usually involves burning fossil fuels. This releases carbon dioxide which causes global warming.

b. Production of Nitric acid (Ostwald’s process)

In the industrial manufacture of nitric acid a catalytic oxidation of ammonia to nitrogen (II) oxide,NO, is carried out then a further oxidation of nitrogen (II) oxide produces nitrogen (IV) oxide, NO2. Nitrogen dioxide is passed through water sprays in a steel absorption tower to produce nitric acid. The excess nitrogen monoxide is recycled back for more oxidation. Platinum is used as a catalyst. There are three steps:

Nitric acid is a strong acid, i.e. dissociates completely when dissolved in water.

i. Chemical properties of nitric acid

Activity 10.3 (b)

Investigating reactions of HNO3with metals

Caution

Work in a well ventilated roomAvoid big yield of the gas when HNO3 reacts with non-metals and metals. The gas is toxic.Concentrated acidis corrosive; much care must be taken in using

Materials and chemicals

Concentrated nitric acid, copper turnings, magnesium ribbon, aluminium powder, carbon (or charcoal), sulphur powder, red and blue litmus paper, spatula, source of heat.

Procedure

1. Collect a pyrex beaker or a clean test tube.

2. Pour 5 cm3 of concentrated HNO3 in the beaker or 2 cm3 of the same acid in the test tube.

3. Add a small piece, 2g to 4g of magnesium to the container.

4. Using wet blue and red litmus paper, check the nature of the gas produced.

5. Record your observations.

6. Repeat the experiment but at step (3) replace magnesium with copper turnings.

7. Record your observations.

This reaction explains why nitric acid is called an “oxidizing acid”; in opposition of acids that cannot dissolve copper called non-oxidizing acids, e.g. HCl(aq), H3PO4(aq).

This reaction can be used to distinguish between a gold sample and a copper sample; gold and copper have almost the same color. If a sample is gold, it does not dissolve in nitric acid; if it is copper, it dissolves in nitric acid to form a blue solution and a brown gas NO2.

Activity 10.3 (c):

Investigate reaction of concentrated HNO3 with non-metals. e.g carbon, and sulphur,

Materials and chemicals

Concentrated nitric acid, carbon (or charcoal), sulphur powder, red and blue litmus paper, spatula, source of heat.

Procedure

1. Have a pyrex beaker.

2. Put one spatula of carbon powder into a beaker; add about 5cm3 of concentrat-ed nitric acid

3. Record your observations.

4. If no reaction is taking place, gently heat the mixture.

5. Test the gas produced with a wet blue and red litmus paper.

6. Record your observations.

Procedure 2

1.Collect a pyrex beaker.

2. Measure one spatula of sulphur powder into the beaker. Add about 5cm3 of concentrated nitric acid.

3. Record your observations.

4. If no reaction is taking place, gently heat the mixture.

5. Test the gas given off with a wet blue and red litmus paper.

6. Record your observations.You can also replace the concentrated HNO3 with dilute nitric acid (1M) and record your observations.

Activity evaluation questions

1. Write balanced equations of the reactions taking place when concentrated nitric acid is used.

2. Write balanced equations of the reactions when dilute nitric acid (1M) is used.

3. Write the formula of the gas given off in all the experiments done.

4. Is the gas released acidic or basic? Explain your answer.

i. Uses of nitric acid

The main industrial use of nitric acid is the production of nitrates as fertilizers. Nitric acid is neutralized with ammonia to give ammonium nitrate.

Others uses are: manufacture of dyes, artificial fiber and drugs, making gun powder and explosives such as trinitrotoluene (TNT)

ii. Environmental impact of industrial production of nitric acid

The manufacture of nitric acid may constitute an environmental problem if appreciable quantities of nitrogen oxides, NOx, are rejected into atmosphere; because they form acid rains, and can contribute the greenhouse effect and global warming.

The main gaseous emissions from the Ostwald process include NO and NO2. Both gases contribute to photochemical smog, and therefore, careful attention must be paid to minimizingthe emission of those gases into the atmosphere. Note that Nitrogen (I) oxide is a significant greenhouse gas.

Checking Up 10.3

1. a. State whether NH3 is a Lewis acid or Lewis base

b. Explain your answer in (a) above

2. Explain briefly how ammonia is prepared in the laboratory.

3. How nitric acid is prepared in the laboratory? Use equation to illustrate your answer.

4. Ammonia is a gas used to manufacture various compounds.

a. State 3 physical properties of ammonia gas.

b. The manufacture of Ammonia in the Haber process is a key industry that is linked to other useful chemicals.

i) Write a balanced equation for the reaction that occurs in the Haber process.

ii) State the conditions of temperature, pressure and catalyst used in this process.

iii. With the help of balanced equations, outline the steps involved in the manufacture of nitric acid from ammonia

5. Write equations to show how ammonia reacts with the following:

a. Oxygen in the absence of a catalyst

b. Oxygen in the presence of a catalyst

c. Copper (II) oxide

d. Hydrochloric acid

6. Write equations to show how nitric acid reacts with the following substances:

a. Copper

b. Sulphur

c. Potassium hydroxide

10.4. Phosphorus and chemical properties of its compounds

Activity 10.4 (a)

Study the figure below and answer the questions asked

1. Explain what is meant by “allotropes” of an element

2. How red phosphorus and white phosphorus differ?

3. Write balanced equations to show how the following oxides are obtained from phosphorous:

a) P2O3

b) P4O104. How does phosphorous react with chlorine?

10.4.1. Allotropes of phosphorus

By definition, allotropy is a property exhibited by some elements to exist in multiple forms with different crystal structures. Allotropes are any two or more physical forms in which an element can exist. Phosphorus exists in two main allotropic forms:

•White phosphorus

•Red phosphorusand

a. White phosphorus (P4)

White phosphorous (P4) exists as molecules made up of four atoms in a tetrahedral structure. Thetetrahedral arrangement results in ring strain and instability. That is why it is very reactive.

When prepared, ordinary phosphorus is white, but it turns light yellow when exposed to sunlight. It is a crystalline, translucent, waxy solid, which glows faintly in moist air and is extremely poisonous.It ignites spontaneously in air at 34°C and must be stored under water. It is insoluble in water, slightly soluble in organic solvents, and very soluble in carbon disulfide. White phosphorus melts at 44.1°C, boils at 280°C. White phosphorus is prepared commercially by heating calcium phosphate with sand (silicon dioxide) and coke in an electric furnace. When heated between 230°C and 300°C in the absence of air, white phosphorus is converted into the red form.White phosphorus spontaneously takes fire in contact with air. White phosphorus is considered and has been used as a chemical weapon.

b. Red phosphorus

It is a micro-crystalline, nonpoisonous powder. It sublimates (passes from the solid state directly to the gaseous state) at 416°C. As shown in the figure below, its structure is made of a chain of P4 units, a kind of polymer: ...(P4)n...

10.4.2. Chemical properties of phosphorous compounds

10.4.3. Phosphoric acid

a. Laboratory preparation of phosphoric acid

In the laboratory phosphoric acid can be prepared by boiling a mixture of red phosphorus with 50% nitric acid in a flask fitted with a reflux condenser on a water bath till no more oxides of nitrogen are liberated. Iodine acts as a catalyst. Phosphoric acid is a weak acid.

b. Reaction of phosphoric acid (H3PO4)with metals and bases

Activity 10.4 (b)

Experiment: Investigating the presence of nitrate ion (NO3-)

Chemicals and apparatus needed are:

1. Freshly prepared iron (II) sulphate solution

2. Concentrated sulphuric acid

3. Test tubes

4. Droppers

Procedure:

1. To about 1cm3 of suspected solution of nitrate, add an equal volume of prepared iron (II) sulphate and tilt the tube.

2. Add cold concentrated H2SO4 slowly on the sides of the test tube.

Study questions

1. Record your observations

2. What is the name given to this test?

10.4.4. Identification of PO43- and NO3- ions

10.4 (b) Identification of NO3- ions

a. Brown ring test: the test carried out in Activity 10.4(b) is called “Brown ring test”. When the test is carried out as indicated, a brown ring is slowly formed between the two layers (acid and aqueous layers).

b. Addition of concentrated sulphuric acid: When a solid nitrate is mixed with about 1cm3 concentrated sulphuric acid and warmed with Cu metal, brown fumes of nitrogen dioxide are formed.

c. Boiling nitrates with powdered Devarda’s alloy (a mixture of copper, zinc and aluminium) with NaOH solution gives a colourless gas with a chocking smell. The gas is ammonia.

Identification of PO43- ions

Formation of immediate yellow precipitate of ammonium phosphomolybdate when a solution of phosphate (V), PO43-, is warmed with nitric acid and ammonium molybdate solution confirms a phosphate (V) ion in solution. Heptaoxodiphosphates (V), P2O74-, and trioxophosphates (III), PO3-3, also give yellow precipitates with the reagents but at a slow rate. The yellow precipitate is due to the hydrolysis of phosphates (V) on warming.

10.4.5. Uses of group 15 elements and compounds

a. Uses of Phosphorus

The main uses of Phosphorus are:

1. Red phosphorus is mainly used on making matches.

2. White phosphorus and zinc phosphate are mainly used as a poison for rats.

3. Phosphorus is used in making phosphor bronze, which is an alloy of copper and tin containing phosphorus.

4. Phosphorus is used for the preparation of phosphorous compounds like P2O5, PCl3,PCl5 and phosphoric acid. Phosphoric acid, in turn, is used in the manufacture of phosphates which are used as fertilizer: NPK (Nitrogen, Phosphate, and Potassi-um) fertilizer.

5. It is used in making incendiary (fire causing) bombs, tracer bullets and for produc-ing smoke screen.

b. Uses of Nitrogen and its compounds

1. Cryopreservation (conservation at low temperature) uses nitrogen to conserve egg, blood, sperm and other biological specimens at very low temperature.

2. Nitrogen gas is used in laboratory or hospital to create inert atmosphere

3. Nitrogen is a constituent of almost every major class of drugs, including antibiot-ics. In the form of nitrous oxide, nitrogen is used as a pharmaceutical anesthetic agent.

4. Ammonia and Nitrates are used as fertilisers to increase soil fertility.

5. Ammonia is used in the manufacture of nitric acid, a starting material for the pro-duction of chemical fertilizers such as ammonium sulphate, ammonium nitrate, and many other kinds of nitrate fertilizers.

6. Ammonia is used in the manufacture of aminobenzene dyes.

7. Nitric acid is used in making explosives such as trinitrotoluene (TNT), and in the production of nitrates.

8. The compound ammonium nitrate mixed with large quantity of salt in water can be used as a freezing mixture.

c. Uses of Arsenic

1. Arsenic is used in semi-conductor electronic devices

2. Arsenic is also used in various agricultural insecticides and poisons.

3. Metallic arsenic forms alloys with lead

Checking Up 10.4

1. What is another name given to yellow phosphorus?

2. Why is white phosphorus stored in water?

3. Write an equation to show how phosphoric acid reacts with calcium hydroxide

4. The chloride and oxide of phosphorus in the higher oxidation state are: PCl5 and P2O5.

a) Give the formulae of the chloride and oxide of phosphorus in a lower oxidation state.

b) Write a balanced equation for the reaction of PCl5 with water.

c) Write a balanced equation for the reaction of P2O5 with water.

d) 25cm3 of the resulting solution of the reaction between P2O5 and water reacted completely with 25cm3 of 0.6moldm-3 NaOH. Calculate the con-centration of the solution.

e) What is the name of the shape of PCl5? Give one of the bond angles in that shape.

10.1.1 Environmental problems of using chemical fertilizers of nitrates and phosphates

Nitric acid is mainly used in the manufacture of nitrates fertilizers. Excess use of nitrates as fertilizers is responsible of one type of pollutions of lakes and rivers called eutrophication.

Eutrophication results from the excessive richness of nutrients in a lake or a water body which causes a dense growth of plant life. When those water plants die and are decomposed, during the decomposition process that uses oxygen, they deplete the oxygen of the water body and render that water incapable of sustaining living aquatic organisms. In that case, the body of water is said to be dead (biologically).

The fraction of the nitrogen-based fertilizers which is not converted to be used by plants accumulates in the soil or lost as run-off. High application rates of nitrogen-containing fertilizers combined with the high water solubility of nitrate leads to increased runoff into surface water as well as leaching into groundwater, thereby causing groundwater pollution.

Ammonium nitrate fertilizers can cause soil acidification:

This may lead to decreases in nutrient availability which may be offset by liming. High levels of fertilizer may cause the breakdown of the symbiotic relationships between plant roots and mycorrhizal fungi.

Conclusion: Use of chemical fertilizers such as nitrates is good because it allows increasing of soil fertility. But we may be careful to avoid that excess of those fertilizers becomes a pollution problems for our soil, lakes and water bodies.

10.5 End unit assessment

Multiple choice questions

1.When heated NH3 is passed over CuO. The gas evolved may be one from the list below. Identify which one.a) N2    b) N2O     c) HNO3     d) NO2

2. When concentrated nitric acid is heated it decomposes to give one from the list below. Identify which one.   a) O2 and N2    b) NO     c) N2O5      d) NO2 and O2

3. Which of the following metal produces nitrous oxide with dilute HNO3

a) Fe     b) Zn     c) Cu      d) Ag

4. Which Nitrogen trihalides is least basic and why?

a) NF3      b) NCl3       c) NBr3       d) NI3

5. P4O6 reacts with water to give one of the following. Identify which one.

a) H3 PO3    b) H4 P2O7       c) HPO3     d) H3 PO4

6. Which one does not form complex as central atom in the list below; justify your answer:

a) N       b) P           c) As         d) Bi

7. Nitrogen dioxide is released by heating one of the following: identify which one.

a) Pb(NO3)2        b) KNO3      c) NaNO2       d) NaNO

Short and long answer Open questions

1. Though nitrogen exhibits +5 oxidation state, it does not form pentahalide. Give a reason.

2. PH3 has lower boiling point than NH3. Why?

3. a. Write the equation for the reaction to prepare ammonia gas from ammonium chloride and calcium hydroxide

b. i) Which other ammonia salt can be used in place of ammonium chloride?

ii) How can you show that ammonia is an alkaline gas?

iii) State three large scale uses of ammonia.

4. a. Describe how nitric acid reacts with copper (your answer should include equations for the reactions and clear observations).

b. Give two reactions that show how nitric acid can behave as an acid and as an oxidizing agent

c. State two uses of HNO3.

5. a. When phosphorous (P) burns in air, it reacts with oxygen to produce an oxide. Write a chemical equation for the reaction of phosphorous with oxygen.

b. Phosphorous pentachloride (PCl5) is one of the chlorides of phosphorous. Deduce the oxidation state of phosphorous atom in phosphorous pentachloride (PCl5) molecule.

c. Give two uses of phosphorous compounds.