18.8 Occurrence, Preparation, and Properties of Phosphorus

Phosphorus: Occurrence, Preparation, and Properties

Phosphorus is one of the most important nonmetals for both biology and industry. It doesn't occur free in nature because it's too reactive, so it must be extracted from mineral sources. This section covers how phosphorus is obtained industrially, the oxides it forms, and how its halides compare to those of nitrogen.

Industrial Preparation of Phosphorus

Phosphorus is obtained from phosphate rock, which contains calcium phosphate $Ca_3(PO_4)_2$. The industrial extraction uses an electric furnace and follows these steps:

1. Phosphate rock is mixed with sand ($SiO_2$) and coke ($C$) in an electric furnace.
2. The mixture is heated to around 1500°C.
3. At that temperature, calcium phosphate reacts with silicon dioxide and carbon:

$2Ca_3(PO_4)_2 + 6SiO_2 + 10C \rightarrow 6CaSiO_3 + P_4 \uparrow + 10CO \uparrow$

1. Phosphorus vapor rises out of the furnace and is condensed under water to prevent it from reacting with oxygen in the air. The product collected this way is white phosphorus.

The byproducts are calcium silicate slag ($CaSiO_3$) and carbon monoxide gas. The coke serves as the reducing agent, pulling oxygen away from the phosphate so elemental phosphorus can form.

Phosphorus Oxides

Phosphorus forms two main oxides, depending on how much oxygen is available during combustion.

Phosphorus(III) oxide $P_4O_6$:

• Forms when white phosphorus burns in a limited oxygen supply (incomplete combustion)
• Reacts with cold water to produce phosphorous acid:

$P_4O_6 + 6H_2O \rightarrow 4H_3PO_3$

• Phosphorous acid ($H_3PO_3$) can act as a reducing agent because it's readily oxidized to phosphoric acid

Phosphorus(V) oxide $P_4O_{10}$:

• Forms when white phosphorus burns in excess oxygen (complete combustion)
• Reacts vigorously with water to produce phosphoric acid:

$P_4O_{10} + 6H_2O \rightarrow 4H_3PO_4$

• Extremely hygroscopic (absorbs moisture from the air), which makes it one of the most effective desiccants available
• Used industrially in the production of phosphoric acid and phosphate fertilizers

The pattern here is straightforward: less oxygen gives the +3 oxide, more oxygen gives the +5 oxide. Each oxide reacts with water to give the corresponding acid.

Phosphorus Halides vs. Nitrogen Halides

Phosphorus halides are significantly more stable than nitrogen halides, and this difference has real practical consequences.

Phosphorus halides:

• $PCl_3$ is a colorless liquid at room temperature; $PCl_5$ is a yellowish-white solid
• Both hydrolyze (react with water), producing the corresponding acid and HCl:

$PCl_3 + 3H_2O \rightarrow H_3PO_3 + 3HCl$

$PCl_5 + 4H_2O \rightarrow H_3PO_4 + 5HCl$

• These compounds are widely used as precursors in making organophosphorus compounds, including pesticides and flame retardants

Nitrogen halides:

• $NCl_3$ is highly unstable and explosive
• It decomposes readily, even at room temperature:

$2NCl_3 \rightarrow N_2 + 3Cl_2$

• Because of this instability, nitrogen halides have very few practical uses

Why the difference? Phosphorus is larger than nitrogen and has a lower electronegativity. Its bigger atomic radius means it can accommodate multiple halogen atoms around it without as much steric strain. Phosphorus can also expand its octet by using its empty 3d orbitals, which is why $PCl_5$ exists but $NCl_5$ does not. Nitrogen is stuck with a maximum of four bonds.

Atomic and Chemical Properties of Phosphorus

• Electron configuration: $[Ne]3s^23p^3$, giving it five valence electrons and a half-filled p subshell
• Oxidation states: -3, +3, and +5 are the most common. The -3 state appears in phosphides (like $PH_3$), while +3 and +5 dominate in oxides and halides
• Electronegativity: Lower than nitrogen (2.19 vs. 3.04 on the Pauling scale), which means P forms less polar bonds with other nonmetals
• Covalent radius: Larger than nitrogen, resulting in longer bond lengths and the ability to form expanded octets using 3d orbitals