P6 P 372b1m

CENIZA, JOHN RAYMUND D. PHYSICS 314 B MWF/TTH 11:00-12:00/3:30-5:00

PHOSPHORUS

1.0 ELEMENT : PHOSPHORUS 2.0 SYMBOL : P 3.0 PHYSICAL PROPERTIES : 3.1 Density/Specific Gravity at 20/4 degrees Celsius: 1.88 (white) (water=1) 2.70 (black) 2.34 (red) 3.2 Boiling Point : 280 degrees Celsius 3.3 Melting Point : 44.1 degrees Celsius (white) 590 degrees Celsius (black/red) 3.4 Vapor Pressure : 1mmHg at 43 atm (white) 3.5 Vapor Density : 4.42 (white) (water=1) 4.77 (red) (water=1) 3.6 Heat of Vaporization: 51.9 kJ/mol (white) 3.7 Heat of Fusion: 0.66 kJ/mol (white) 3.8 Molar Heat Capacity: 23.824 J/(mol-K) 3.9 Molecular Weight: 30.973762 3.10 Molar volume: 196 cm3 4.0 CHEMICAL PROPERTIES: 4.1 Reaction of phosphorus with oxygen:

White phosphorus glows in the dark when exposed to damp air in a process known as chemiluminescence. White phosphorus must be handled with great care. It spontaneously ignites in air at about room temperature to form "phosphorus pentoxide" - actually tetraphosphorus decaoxide, P4O10. P4(s) + 5O2 (g) → P4O10(s) 4.2 Under careful control (75% O2, 25% N2, 50°C, 90 mm Hg), a mixture is formed, one of the products within which is “phosphorus trioxide” actually tetraphosphorus hexaoxide P4O6. P4(s) + 3O2 (g) → P4O6(s) PHOSPHORUS REACTS WITH THE HALOGENS White phosphorus, P4, reacts vigorously with all the halogens at room temperature to form phosphorus trihalides. 4.3 Reaction of phosphorus to fluorine: It reacts with fluorine, F2. P4(s) + 6F2 (g) → 4PF3(g)

4.4 Reaction of phosphorus to chlorine:

It reacts with chlorine, Cl2. White phosphorus burns spontaneously in chlorine to produce a mixture of two chlorides, phosphorus (III) chloride and phosphorus (V) chloride (phosphorus trichloride and phosphorus pentachloride). P4(s) + 6Cl2(g) → 4PCl3 (l) 4.5 Reaction of phosphorus to bromine: It reacts with bromine, Br2. P4(s) + 6Br2 (g) → 4PBr3 (l) 4.6 Reaction of phosphorus to iodine: It reacts with iodine, I2. P4(s) + 6I2 (g) → 4PI3 (g) 4.7 Phosphorus is a white solid with a waxy appearance, which melts at 44.1 oC and boils at 287oC. It is made by reducing calcium phosphate with carbon in the presence of silica (sand) at very high temperatures. 2 Ca3(PO4)2(s) + 6 SiO2(s) + 10 C(s) 6 CaSiO3(s) + P4(s) + 10 CO(g) 4.8 P4O6 and P4O10 react with water: P4O6 and P4O10 react with water to form phosphoric acid, H 3PO4. P4O10(s) + 6 H2O(l) 4 H3PO4(aq) 4.9 P4O6 and P4O10 react with water: P4O6 and P4O10 react with water to form phosphorous acid, H 3PO3. P4O6(s) + 6 H2O(l) 4 H3PO3(aq) 4.10 Metal phosphides react with water: These metal phosphides react with water to produce a poisonous, highly reactive, colorless gas known as phosphine (PH 3), which has the foulest odor the authors have encountered. Ca3P2(s) + 6 H2O(l) 2 PH3(g) + 3 Ca2+(aq) + 6 OH-(aq)

5.0 HISTORY: In 1669, a German merchant called Henning Brand obtained elementary phosphorus through the distillation of urine, writing a letter to Leibniz reporting its discovery. It is quite probable that, in the 12th century, Arabian alchemists have obtained the element using this process. However, the credit is given to Brand. The name of phosphorus has a Greek origin meaning "it possesses brilliance" due to its property of shining in the darkness when exposed to the air. About one century after its original work, Brand discovered that phosphorus is an important constituent of the bones, introducing a new method of industrial production of phosphorus. The reaction of the bones with nitric or sulfuric acid produces phosphoric acid that, when heated up with coal, produces elementary phosphorus. This was the first method of phosphorus commercial production. 6.0 APPLICATIONS TO PHYSICS: 6.1 Used in the manufacture of gasoline additives: Phosphorus oxychloride (POCl3) is a compound of phosphorus with a number of uses like in the manufacturing of gasoline additives, fire retardant agent, manufacture of transistors for electronic devices. 6.2 Used to make fertilizers: In 1996, 91% of all the phosphate rock mined in the United States was used to make fertilizers. 6.3 Important for energy transfer in cells: Phosphorus is essential to all living things. It forms the sugar phosphate backbone of DNA and RNA. 6.4 Ingredients in some detergents: They are beginning to be phased out in some countries because they can lead to high phosphate levels in natural water supplies causing unwanted algae to grow. 6.5 Used to make safety matches: The phosphorus pentasulfide or phosphorus sequisulfide coat the tip of the match and bursts into flame. It ignites other chemicals on the head of the match. 6.6 Useful in medical studies: Radioactive phosphorus can be used as a tracer to study parts of the body. It can also determine how much blood is in a person’s body and it can also help locate the presence of tumors in the brain, eyes, breasts and skin.

SARSALEJO, JOFFRY PHYSICS 314 C MWF/TTH 2:00-3:00/5:00-6:30

ANTIMONY

1.0 ELEMENT : Antimony

2.0 SYMBOL : Sb 3.0 PHYSICAL PROPERTIES : 3.1 Molecular Weight: 30.973762 3.2 Boiling Point : 1908 K 3.3 Melting Point : 903.78 K 3.4 Vapor Pressure : 1mmHg at 43 atm (white) 3.5 Vapor Density: 4.42) 3.6 Density: 6.697 g/cm^3 3.7 Molar volume: 196 cm3 3.8 Molar Heat Capacity: 25.23 J/(mol-K) 3.9 Heat of Vaporization: 193.43 kJ/mol 3.10 Heat of Fusion: 19.79 kJ/mol 4.0 CHEMICAL PROPERTIES: 4.1 Reaction of antimony with acids: Antimony dissolves in hot concentrated sulphuric acid, H2SO4, or nitric acid, HNO3, to form solutions containing Sb(III). The sulphuric acid reaction produces sulphur(IV) dioxide gas. Antimony does not react with hydrochloric acid in the absence of oxygen. 4.2 Reaction of antimony with oxygen: Upon heating, antimony reacts with oxygen in air to form the trioxide antimony (III) oxide, Sb2O3. The flame is bluish white.

4Sb(s) + 3O2 (g) → 2sb2O3(s) 4.3 Reaction of antimony with water: At red heat, antimony reacts with water to form the trioxide antimony (III) oxide, Sb2O3. Antimony reacts more slowly at ambient temperatures.

2Sb(s) + 3H2O (g) → Sb2O3(s) + 3H2 (g) ANTIMONY REACTS WITH THE HALOGENS Antimony reacts under controlled conditions with the halogens to form the respective trihalides antimony (III)

4.4 Reaction of antimony to fluorine: It reacts with fluorine, F2. 2Sb(s) + 3F2 (g) → 2SbF3(s) [white] 4.5 Reaction of antimony to chlorine: It reacts with chlorine, Cl2. 2Sb(s) + 3Cl2(g) → 2SbCl3(s) [white]

4.6 Reaction of antimony to iodine: It reacts with iodine, I2. 2Sb(s) + 3I2 (g) → 2SbI3(s) [red] 4.7 Reaction of antimony to bromine: It reacts with bromine, B2. 2Sb(s) + 3Br2 (g) → 2SbBr3(s) [white] 4.8 Antimonides reacts with acids: These antimonides with acid produces the unstable gas stibine, SbH3.

Sb3− + 3 H+ → Sb3 4.9 Reaction of antimony to chlorine: Antimony Pentachloride, SbCl5 was first obtained by Rose by the direct union of the elements, powdered antimony combining spontaneously with chlorine, with incandescence. No antimony trichloride is formed. It is usually prepared by saturating molten antimony trichloride with chlorine, and distilling under reduced pressure. 4.10 Reaction with 15M HNO3: (mild, =>Sb2O5)

5.0 HISTORY: Antimony was recognized in predynastic Egypt as an eye cosmetic as early as about 3100 B.C when the cosmetic palette was invented. An artifact said to be a part of a vase, made of antimony dating to about 3000 B.C was found at Telloh, Chaldea (part present Day Iraq) and a copper object plated with antimony dating between 2500 B.B and 2200 B.C has been found in Egypt. The element was probably first named by Roman scholar Pliny (AMP 2379) who called it stibium. Muslim chemist Abu Musa Jabir IbnHaygan probably first called it antimony anti (“not”) and monos (“alone”). The name comesfrom the fact that antimony does not occur alone in nature. The first detailed reports about antimony were published in 1707 when French chemist Nicolas Lemory (1645-1715) published his famous book, Treatise on Antimony.

6.0 APPLICATIONS TO PHYSICS: 6.1 Flame retardants: Antimony is mainly used as it’s trioxide in making flame proofing compounds. It is used in combination with halogenated flame retardants with the only exception being in halogen containing polymers. 6.2 Lead Acid Batteries: Antimony forms a highly useful alloy with lead and added on batteries thus improves the charging characteristics and reduces generation of unwanted hydrogen during charging. 6.3 Fining agent for TV screens: Antimony removes microscopic bubbles in glass. This is achieved by the interaction of antimony ions with oxygen, interfering the latter from forming bubbles. 6.4 Used in Semiconductor Industry: Antimony is used as a dopant for heavily doped n-type silicon wafers and in the production of diodes, infrared detectors and Hall effect devices. 6.5 Veterinary Preparations:

Antimony and its compounds are used in several veterinary preparations like anthiomaline or lithium antimony thiomolate which is used as a skin conditioner in rumnants. Antimony has a nourishing or conditioning effect on keratinized tissues at least in animals.

CORTES, CHAD PHYSICS 314 C MWF/TTH 2:00-3:00/5:00-6:30

PHOSPHORUS

4.0 ELEMENT : PHOSPHORUS 5.0 SYMBOL : P 6.0 PHYSICAL PROPERTIES : 6.1 Molecular Weight : 30.973762 6.2 Boiling Point : 280 degrees Celsius 6.3 Melting Point : 44.1 degrees Celsius (white) 590 degrees Celsius (black/red) 6.4 Vapor Pressure : 1mmHg at 43 atm (white) 6.5 Vapor Density : 4.42 (white) (water=1) 4.77 (red) (water=1) 3.6 Density/Specific Gravity at 20/4 degrees Celsius: 1.88 (white) (water=1) 2.70 (black) 2.34 (red) 3.7 Molar volume: 196 cm3 3.8 Molar Heat Capacity: 23.824 J/(mol-K) 3.9 Heat of Vaporization: 51.9 kJ/mol (white) 3.10 Heat of Fusion: 0.66 kJ/mol (white) 4.0 CHEMICAL PROPERTIES:

4.1 Reaction of phosphorus with oxygen:

White phosphorus glows in the dark when exposed to damp air in a process known as chemiluminescence. White phosphorus must be handled with great care. It spontaneously ignites in air at about room temperature to form "phosphorus pentoxide" - actually tetraphosphorus decaoxide, P4O10. P4(s) + 5O2 (g) → P4O10(s) 4.2 Under careful control (75% O2, 25% N2, 50°C, 90 mm Hg), a mixture is formed, one of the products within which is “phosphorus trioxide” actually tetraphosphorus hexaoxide P4O6. P4(s) + 3O2 (g) → P4O6(s) PHOSPHORUS REACTS WITH THE HALOGENS White phosphorus, P4, reacts vigorously with all the halogens at room temperature to form phosphorus trihalides. 4.3 Reaction of phosphorus to fluorine: It reacts with fluorine, F2. P4(s) + 6F2 (g) → 4PF3(g)

4.4 Reaction of phosphorus to chlorine:

It reacts with chlorine, Cl2. White phosphorus burns spontaneously in chlorine to produce a mixture of two chlorides, phosphorus (III) chloride and phosphorus (V) chloride (phosphorus trichloride and phosphorus pentachloride). P4(s) + 6Cl2(g) → 4PCl3 (l) 4.5 Reaction of phosphorus to bromine: It reacts with bromine, Br2. P4(s) + 6Br2 (g) → 4PBr3 (l) 4.6 Reaction of phosphorus to iodine: It reacts with iodine, I2. P4(s) + 6I2 (g) → 4PI3 (g) 4.7 Phosphorus is a white solid with a waxy appearance, which melts at 44.1 oC and boils at 287oC. It is made by reducing calcium phosphate with carbon in the presence of silica (sand) at very high temperatures. 2 Ca3(PO4)2(s) + 6 SiO2(s) + 10 C(s) 6 CaSiO3(s) + P4(s) + 10 CO(g) 4.8 P4O6 and P4O10 react with water: P4O6 and P4O10 react with water to form phosphoric acid, H 3PO4.

P4O10(s)

+

6 H2O(l)

4 H3PO4(aq)

4.9 P4O6 and P4O10 react with water: P4O6 and P4O10 react with water to form phosphorous acid, H 3PO3. P4O6(s) + 6 H2O(l) 4 H3PO3(aq) 4.10 Metal phosphides react with water: These metal phosphides react with water to produce a poisonous, highly reactive, colorless gas known as phosphine (PH 3), which has the foulest odor the authors have encountered. Ca3P2(s) + 6 H2O(l) 2 PH3(g) + 3 Ca2+(aq) + 6 OH-(aq)

6.0 HISTORY: Phosphorus came from the Greek word “phosphoros” meaning “bringer of light” was discovered in 1669 at by a German merchant called Henning Brand obtained elementary phosphorus through the distillation of urine and writing a letter to Leibniz reporting its discovery. About one century after its original work, Brand discovered that phosphorus is an important constituent of the bones thus introducing a new method of industrial production of phosphorus. At the end of 18th century, James Rean developed the first process for the production of the element with an electric furnace. Rean’s method to obtain elementary phosphorus remains in present technology. 6.0 APPLICATIONS TO PHYSICS: 6.1 Useful in medical studies: Radioactive phosphorus can be used as a tracer to study parts of the body. It can also determine how much blood is in a person’s body and it can also help locate the presence of tumors in the brain, eyes, breasts and skin. 6.2 Used to make fertilizers: In 1996, 91% of all the phosphate rock mined in the United States was used to make fertilizers. 6.3 Used to make safety matches: The phosphorus pentasulfide or phosphorus sequisulfide coat the tip of the match and bursts into flame. It ignites other chemicals on the head of the match. 6.4 Ingredients in some detergents: They are beginning to be phased out in some countries because they can lead to high phosphate levels in natural water supplies causing

unwanted algae to grow. 6.5 Important for energy transfer in cells: Phosphorus is essential to all living things. It forms the sugar phosphate backbone of DNA and RNA. 6.6 Used in the manufacture of gasoline additives: Phosphorus oxychloride (POCl3) is a compound of phosphorus with a number of uses like in the manufacturing of gasoline additives, fire retardant agent, manufacture of transistors for electronic devices.