Nitrogen and Phosphorus

Properties of Nitrogen

Nitrogen in its elemental form is a non-metallic gas that makes up 78 percent of Earth’s atmosphere.

Learning Objectives

Discuss the properties of nitrogen.

Key Takeaways

Key Points

  • Nitrogen is a chemical element with symbol N and atomic number 7. Elemental nitrogen is a colorless, odorless, tasteless, and mostly inert diatomic gas at standard conditions, constituting 78.09 percent of Earth’s atmosphere by volume.
  • Nitrogen gas is an industrial gas produced by the fractional distillation of liquid air or by mechanical means using gaseous air. Commercial nitrogen is often a byproduct of air-processing for industrial concentration of oxygen.
  • Nitrogen gas has a variety of applications, including serving as an inert replacement for air where oxidation is undesirable. Liquid nitrogen is also used to cryogenically freeze objects.

Key Terms

  • nitrogen: A chemical element (symbol N) with an atomic number of 7 and atomic weight of 14.0067 amu.
  • amino acid: Generally, molecules that contain both an amino and a carboxylic acid functional group. The monomers from which polypeptide chains, or proteins, are built are amino acids.
  • elemental: Of, relating to, or being an element (as opposed to a compound).

The element nitrogen was discovered as a separable component of air by Scottish physician Daniel Rutherford in 1772. Nitrogen compounds were well known during the Middle Ages. Alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpetre (sodium nitrate or potassium nitrate), most notably in gunpowder and later as fertilizer.

Nitrogen is a chemical element with symbol N and atomic number 7. Elemental nitrogen is a colorless, odorless, tasteless, and mostly inert diatomic gas at standard conditions, constituting 78.09 percent of Earth’s atmosphere by volume. Nitrogen is a common element in the universe, estimated at about seventh in total abundance in our galaxy and the solar system. Its occurrence there is thought to be entirely due to synthesis by fusion of carbon and hydrogen in supernovas. Due to the volatility of elemental nitrogen and its compounds with hydrogen and oxygen, nitrogen is far less common on the rocky planets of the inner solar system and is a relatively rare element on Earth. However, as on Earth, nitrogen and its compounds occur commonly as gases in the atmospheres of planets and moons.

Nitrogen in Living Systems

Nitrogen occurs in all living organisms, primarily in amino acids which make up proteins, and nucleic acids (DNA and RNA). The human body is about three percent nitrogen by weight, the fourth-most abundant element after oxygen, carbon, and hydrogen. Nitrogen resides in the chemical structure of almost all neurotransmitters and is a defining component of alkaloids, biological molecules produced as secondary metabolites by many organisms.

The nitrogen cycle describes the movement of the element from the air into the biosphere and organic compounds and back into the atmosphere. Synthetically produced nitrates are key ingredients of industrial fertilizers and key pollutants causing the eutrophication of water systems.

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The Nitrogen Cycle: The figure summarizes the major processes through which nitrogen is converted between its various forms on the surface of the earth.

Industrial Production of Nitrogen

Nitrogen gas is an industrial gas produced by the fractional distillation of liquid air or by mechanical means using gaseous air (i.e., pressurized reverse osmosis membrane or pressure swing adsorption). Commercial nitrogen is often a byproduct of air-processing for industrial concentration of oxygen for steelmaking and other purposes. When supplied compressed in cylinders, it is often called OFN (oxygen-free nitrogen).

In a chemical laboratory it is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite, or through the decomposition of sodium azide:

[latex]\text{NH}_4\text{Cl} (\text{aq}) + \text{NaNO}_2 (\text{aq}) \rightarrow \text{N}_2 (\text{g}) + \text{NaCl} (\text{aq}) + 2 \text{H}_2\text{O} (\text{l})[/latex]
[latex]2 \text{NaN}_3 \rightarrow 2 \text{Na} + 3 \text{N}_2[/latex]

Chemical Properties of Nitrogen

Nitrogen is a nonmetal with an electronegativity of 3.04. It has five electrons in its outer shell and is, therefore, trivalent in most compounds. The triple bond in molecular nitrogen (N2) is one of the strongest known. The resulting difficulty of converting N2 into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N2, have dominated the role of nitrogen in both nature and human economic activities.

Nitrogen Emission Spectrum

Molecular nitrogen (14N2) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and, therefore, has no dipole moment to couple the electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths, beginning at a wavelength of around 100 nanometers. This is associated with electronic transitions in the molecule to states in which charge is not distributed evenly between nitrogen atoms. Nitrogen absorption leads to significant absorption of ultraviolet radiation in the Earth’s upper atmosphere and the atmospheres of other planetary bodies.

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Spectrum of Nitrogen: Sending an electric current through nitrogen excites the electrons to higher energy levels. When they fall to lower energy levels, certain frequencies of light (based on the difference in energy of the energy levels) are observed, as shown.

Applications of Nitrogen Gas

Nitrogen gas has a variety of applications, including:

  • As an inert replacement for air where oxidation is undesirable
  • As a modified atmosphere, pure or mixed with carbon dioxide, to preserve the freshness of packaged or bulk foods
  • In ordinary incandescent light bulbs as an inexpensive alternative to argon
  • In production of electronic parts such as transistors, diodes, and integrated circuits
  • Filling automotive and aircraft tires due to its inertness and lack of moisture or oxidative qualities, as compared to air
  • As a propellant for draft wine, and as an alternative to or in combination with carbon dioxide in carbonated beverages

Nitrogen is also used in preparing samples for chemical analysis to concentrate and reduce the volume of liquid samples. Directing a pressurized stream of nitrogen gas perpendicular to the surface of the liquid allows the solvent to evaporate while leaving the solute(s) and unevaporated solvent behind. Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball guns. But, nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.

Nitrogen Compounds

Nitrogen compounds, and especially their oxidized derivatives, are important in biological systems and as explosives.

Learning Objectives

Give examples of applications of nitrogen compounds.

Key Takeaways

Key Points

  • Nitrogen oxides, called NOx compounds, are important for their explosive properties. These properties are determined by the extremely strong and stable bond found in molecular, diatomic nitrogen, N2, which has a bond dissociation energy of 945 kJ/mol (226 kcal/mol).
  • The main neutral hydride of nitrogen is ammonia (NH3), which has a pKb of 9.2, and is thus a weak base. The corresponding deprotonated species, NH2, is called an amide and is a strong base (because it’s the conjugate base of ammonia, whose pKa is around 38).
  • Nitrogen is a constituent of molecules in every major drug class in pharmacology and medicine, from antibiotics to neurotransmitters and beyond. One important aspect of nitrogen is that it is the only non-metal that can maintain a positive charge at physiological pH.

Key Terms

  • propellant: Fuel, oxidizer, reaction mass or mixture for one or more engines (especially internal combustion engines or jet engines) that is carried within a vehicle prior to use.
  • oxide: A binary chemical compound of oxygen with another chemical element.
  • anion: A negatively charged ion, as opposed to a cation.
  • bond dissociation energy: The energy required to separate two atoms joined by a particular bond. Expressed in terms of a mole of such bonded atoms. Indicates the strength of the bond.

Survey of Nitrogen Compounds and their Uses

Nitrogen compounds play an important role in many aspects of life and commercial processes, from the industrial production of fertilizers to the building blocks of life.

The nitrogen-nitrogen triple bond in N2 contains 226 kcal/mol of energy, making it one of the strongest bonds known. When nitrogen gas is formed as a product from various reactions, the bond energy associated with the N-N triple bond is released, causing the explosive properties seen in many nitrogen compounds.

Amines

The main neutral hydride of nitrogen is ammonia (NH3), although hydrazine (N2H4) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution, ammonia forms the ammonium ion (NH4+). The pKa of ammonium chloride is 9.2. Liquid ammonia (boiling point 240 K) is amphiprotic (displaying either Brønsted-Lowry acidic or basic character) and forms ammonium and the less common amide ions (NH2). Ammonia has a pKa of 38, making the corresponding amide ions very strong bases. Singly, doubly, triply and quadruply substituted alkyl compounds of ammonia are called amines (four substitutions, to form commercially and biologically important quaternary amines, results in a positively charged nitrogen, and thus a water-soluble compound).

Azides

Other classes of nitrogen anions (negatively charged ions) are the poisonous azides (N3), which are linear and isoelectronic to carbon dioxide, but which bind to important iron-containing enzymes in the body in a manner resembling cyanide.

Nitrogen Oxides

Another molecule of the same structure is the colorless and relatively inert anesthetic gas nitrous oxide (dinitrogen monoxide, N2O), also known as laughing gas. This is one of a variety of nitrogen oxides that form a family often abbreviated as NOx. Nitric oxide (nitrogen monoxide, NO), is a natural free radical used in signal transduction in both plants and animals. The reddish and poisonous nitrogen dioxide (NO2) contains an unpaired electron and is an important component of smog. Nitrogen molecules containing unpaired electrons show a tendency to dimerize (thus pairing the electrons), and are, in general, highly reactive. The corresponding acids are nitrous (HNO2) and nitric acid (HNO3), with the corresponding salts called nitrites and nitrates.

Nitrogen Compounds used as Explosives and Propellants

One of the earliest uses of a nitrogen compound as an explosive was potassium nitrate, also called saltpeter, used in gunpowder. This is a mixture of potassium nitrate, carbon and sulfur. When the mixture is ignited in an enclosed space, such as a gun-barrel or a firework, the nitrate ions oxidize the carbon and sulfur in a highly exothermic reaction, producing high- temperature gases very rapidly. This can propel a bullet out of a gun or cause a firework to explode.

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Potassium Nitrate: Potassium nitrate, or saltpeter, was a constituent of early gunpowder. It serves as an oxidant, along with an oxidizable material such as sugar.

The higher oxides, dinitrogen trioxide (N2O3), dinitrogen tetroxide (N2O4) and dinitrogen pentoxide (N2O5), are unstable and explosive, a consequence of the chemical stability of N2. Nearly every hypergolic (i.e. not requiring ignition) rocket engine uses N2O4 as the oxidizer; their fuels, various forms of hydrazine, are also nitrogen compounds.

These engines were extensively used on spacecraft such as the space shuttle and those of the Apollo Program because their propellants are liquids at room temperature and ignition occurs on contact without an ignition system, allowing many precisely controlled burns. N2O4 is an intermediate in the manufacture of nitric acid HNO3, one of the few acids stronger than the hydronium ion, and a fairly strong oxidizing agent.

Nitrogen is notable for the range of explosively unstable compounds that it can produce. Nitrogen triiodide (NI3) is an extremely sensitive contact explosive. Nitrocellulose, produced by nitration of cellulose with nitric acid, is also known as guncotton. Nitroglycerin, made by nitration of glycerin, is the dangerously unstable explosive ingredient of dynamite. The comparatively stable, but less powerful explosive trinitrotoluene (TNT) is the standard explosive against which the power of nuclear explosions are measured. In all cases, the explosive properties of nitrogen compounds are derived from the extreme stability of the product of these reactions: gaseous molecular nitrogen, N2.

Nitrogen Compounds in Drugs and Medicine

Nitrogen is a constituent of molecules in every major drug class in pharmacology and medicine. Nitric oxide (NO) has recently been discovered to be an important signaling molecule in physiology. Nitrous oxide (N2O) was discovered early in the 19th century to be a partial anesthetic, though it was not used as a surgical anesthetic until later. Called “laughing gas,” it was found to induce a state of social disinhibition resembling drunkenness.

Other notable nitrogen-containing drugs are drugs derived from plant alkaloids, such as morphine. Many alkaloids are known to have pharmacological effects; in some cases, they appear as natural chemical defenses of plants against predation. Drugs that contain nitrogen include all major classes of antibiotics, and organic nitrate drugs like nitroglycerin and nitroprusside that regulate blood pressure and heart action. Amines (alkyl derivatives of nitrogen) are important in pharmacology because they can readily carry a positive charge, as the corresponding protonated ammonium species. This allows for electrostatic interactions between the ammonium cation and various negatively charged or polarizable species in proteins.

Properties of Phosphorus

Phosphorus is found in its elemental form as different allotropes, none of which are stable in the presence of oxygen.

Learning Objectives

Review the properties of phosphorus.

Key Takeaways

Key Points

  • Phosphorus is a chemical element with symbol P and atomic number 15. A multivalent nonmetal of the nitrogen group, phosphorus as a mineral is almost always present in its maximally oxidized state, as inorganic phosphate.
  • Phosphorus is essential for life. As part of the phosphate group, it is a component of DNA, RNA, ATP (adenosine triphosphate), and the phospholipids that form all cell membranes.
  • Phosphorus exists in several forms ( allotropes ) that exhibit strikingly different properties. The two most common allotropes are white phosphorus and red phosphorus.

Key Terms

  • phosphate: Any salt or ester of phosphoric acid
  • adenosine triphosphate: A nucleotide that occurs in biological organisms and is used as a source of energy in cellular reactions and processes.
  • allotrope: Any form of a pure element that has a distinctly different molecular structure to another form of the same element.

Phosphorus is a chemical element with symbol P and atomic number 15. A multivalent nonmetal of the nitrogen group, phosphorus as a mineral is almost always present in its maximally oxidized state, as inorganic phosphate rocks. Elemental phosphorus exists in two major forms — white phosphorus and red phosphorus — but due to its high reactivity, phosphorus is never found as a free element on Earth.

While the term “phosphorescence” is derived from the ability of white phosphorus to glow faintly upon exposure to oxygen, the current chemical understanding is that this phenomenon is actually chemiluminescence, a mechanism of light emission distinct from phosphorescence.

Importance of Phosphorus

Phosphorus is essential for life. As evidence of the link between phosphorus and terrestrial life, elemental phosphorus was historically first isolated from human urine, and bone ash was an important early phosphate source. As phosphate, it is a component of DNA, RNA, ATP (adenosine triphosphate), and the phospholipids that form all cell membranes. Low phosphate levels are an important limit to growth in some aquatic systems, and the chief commercial use of phosphorus compounds for production of fertilizers is due to the need to replace the phosphorus that plants remove from the soil.

Phosphorus exists in several forms (allotropes) that exhibit strikingly different properties.

  • The two most common allotropes are white phosphorus and red phosphorus.
  • Another form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight.
  • Black phosphorus is obtained by heating white phosphorus under high pressures (about 12,000 standard atmospheres, or 1.2 gigapascals). In appearance, properties, and structure, black phosphorus resembles graphite — it is black and flaky, a conductor of electricity, and has puckered sheets of linked atoms.
  • Another allotrope is diphosphorus; it contains a phosphorus dimer as a structural unit and is highly reactive.

White Phosphorus and Related Molecular Forms

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Diagram of White Phosphorus: The most important elemental form of phosphorus is white phosphorus, P4, which exhibits the bonding shown.

The most important elemental form of phosphorus in terms of applications is white phosphorus. It consists of tetrahedral P4 molecules, in which each atom is bound to the other three atoms by a single bond. This P4 tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C, at which point it starts decomposing into P2 molecules. Solid white phosphorus exists in two forms; at low temperatures, the β form is stable, and at high temperatures, the α form is predominant. These forms differ in terms of the relative orientations of the constituent P4 tetrahedra.

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White Phosphorus: White phosphorus must be stored in an inert medium away from oxygen, such as in mineral oil, as shown here.

White phosphorus is the least stable, the most reactive, the most volatile, the least dense, and the most toxic of the allotropes. It gradually changes to red phosphorus, a transformation accelerated by light and heat. Samples of white phosphorus almost always contain some red phosphorus and so appear yellow. For this reason it is also called yellow phosphorus. It glows in the dark (when exposed to oxygen) with a very faint tinge of green and blue, and it is highly flammable and pyrophoric (self-igniting) upon contact with air. It is also toxic, causing severe liver damage upon ingestion. Owing to its pyrophoricity, white phosphorus is used as an additive in napalm. The odor of combustion of this form has a characteristic garlic smell, and samples are commonly coated with white “(di)phosphorus pentoxide,” which consists of P4O10 tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is insoluble in water but soluble in carbon disulfide.

Red Phosphorus

Red phosphorus is polymeric in structure. It can be viewed as a derivative of P4 — one of the P-P bonds is broken, and one additional bond is formed between the neighboring tetrahedrons, resulting in a chain-like structure. Red phosphorus may be formed by heating white phosphorus to 250 °C (482 °F) or by exposing it to sunlight. Phosphorus after this treatment is amorphous. Upon further heating, this material crystallizes. In this sense, red phosphorus is not an allotrope, but rather an intermediate phase between white and violet phosphorus, and most of its properties have a range of values. For example, freshly prepared, bright-red phosphorus is highly reactive and ignites at about 300 °C, though it is still more stable than white phosphorus, which ignites at about 30 °C. After prolonged heating or storage, the color darkens; the resulting product is more stable and does not spontaneously ignite in air.

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Red Phosphorus Crystal Structure: Red phosphorus is similar to P4 but is polymeric: one of the P-P bonds has been broken and is now attached to the next P4 unit.

Phosphorus Production

About 1,000,000 short tons (910,000 t) of elemental phosphorus is produced annually. Calcium phosphate (phosphate rock), mostly mined in Florida and North Africa, can be heated to 1,200-1,500 °C with sand, which is mostly SiO2, and coke (impure carbon) to produce vaporized P4. The product is subsequently condensed into a white powder underwater to prevent oxidation by air.

Phosphorus Compounds

Phosphorus compounds consist mostly of compounds containing strong phosphorus-oxygen bonds. They are important in fertilizers and biology.

Learning Objectives

Discuss the chemistry and the biological importance of phosphorous compounds.

Key Takeaways

Key Points

  • The chemistry of phosphorus is often dominated by the strength of the oxygen -phosphorus bond, which is around 152 kcal/mol. This extremely strong bond is the driving force behind several reactions involving phosphorus.
  • The majority of phosphorus-containing compounds are produced for use as fertilizers. For this purpose, phosphate -containing minerals are converted to phosphoric acid.
  • Inorganic phosphorus in the form of the phosphate PO43− is required for all known forms of life. It plays a major role in biological molecules; for example, it forms part of the structural framework of DNA and RNA.

Key Terms

  • kilocalorie: A non-SI unit of energy equal to 1,000 calories, used (now rarely) in chemistry or physics; equal to 1 calorie or Calorie as used in nutrition (symbol: kcal).

The Chemistry of Phosphorus Compounds

The chemistry of phosphorus is often dominated by the strength of the oxygen-phosphorus bond, which is around 152 kcal/mol (kilocalories per mole). This extremely strong bond is the driving force behind several reactions involving phosphorus. For example, the reaction of PCl5 with water to become H3PO4 allows it to serve as a drying agent, or dessicant, with the P-O bond formation as the driving force. The oxygen-phosphorus bond also prohibits phosphorus from being observed in its elemental state in nature. It is always found as an oxide.

The majority of phosphorus-containing compounds are produced for use as fertilizers. For this purpose, phosphate-containing minerals are converted to phosphoric acid. Two distinct routes are employed; the main one is treatment of phosphate minerals with sulfuric acid. The other process utilizes white phosphorus, which may be produced by reaction and distillation from very low-grade phosphate sources. The white phosphorus is then oxidized to phosphoric acid and finally neutralized with a base to yield phosphate salts. Phosphoric acid obtained from white phosphorus is relatively pure and is the main source of phosphates used in detergents and other non-fertilizer applications.

Biological Significance

Inorganic phosphorus in the form of the phosphate PO43− is required for all known forms of life. It plays a major role in biological molecules; for example, it forms part of the structural framework of DNA and RNA. As such, phosphate salts are used as fertilizers to aid plant growth. Living cells also use phosphate to transport cellular energy in the form of adenosine triphosphate (ATP). Nearly every cellular process that uses energy obtains it in the form of ATP. ATP is also important for phosphorylation, a key regulatory and signal-transducing event in cells. Phospholipids are the main structural components of all cellular membranes and consist of a long alkyl chain terminating in a phosphate group. Calcium phosphate salts help to harden bones.

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Phosphate in DNA: DNA strands have a phosphate-deoxyribose backbone. Two DNA strands are shown in the figure.

Living cells are defined by a membrane that separates it from its surroundings. Biological membranes are made from a phospholipid matrix and proteins, typically in the form of a bilayer. Phospholipids are derived from glycerol, such that two of the glycerol hydroxyl (OH) groups have been replaced with fatty acids to form ester linkages, and the third hydroxyl group has been replaced with a phosphate group.

Oxoacids of Phosphorus

Phosphorous oxoacids are extensive, often commercially important, and sometimes structurally complicated. They all have acidic protons bound to oxygen atoms, and some have nonacidic protons that are bonded directly to phosphorus. Although many oxoacids of phosphorus are formed, only nine are important, and three are crucial: hypophosphorous acid, phosphorous acid, and phosphoric acid.

Phosphorus with an oxidation state of +1:

  • Hypophosphorous acid, H3PO2, contains one acidic OH bond and two (relatively) non-acidic PH bonds.

Phosphorus with an oxidation state of +3:

  • Phosphorous acid, H3PO3, contains two acidic OH bonds and one PH bond.
  • Orthophosphorous acid, also written H3PO3, contains three acidic OH bonds and no PH bonds.

Phosphorus with an oxidation state of +5:

  • Orthophosphoric acid, H3PO4, is the parent acid and most common oxidation state of phosphorus, with three acidic OH protons. Condensation between two phosphoric acid groups can lead to polyphosphates, such as meta- and polyphosphoric acid.
  • Metaphosphoric acid, (HPO3)n, which occurs when phorphoric-acid molecules become bound together in ring structures, each vertex of the ring containing one acidic OH proton.
  • Polyphosphoric acid, H(HPO3)nOH, which consists of multiple orthophosphoric acids bound together, each via a common oxygen.

Organophosphorus Compounds

Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The PCl3 serves as a source of P+3 in routes to organophosphorus (III) compounds. For example, it is the precursor to triphenylphosphine:

[latex]\text{PCl}_3 + 6\text{Na} + 3\text{C}_6\text{H}_5\text{Cl} \rightarrow \text{P}(\text{C}_6\text{H}_5)_3 + 6\text{NaC}[/latex]

Treatment of phosphorus trihalides with alcohols and phenols yields phosphites, such as triphenylphosphite:

[latex]\text{PCl}_3 + 3\text{C}_6\text{H}_5\text{OH} \rightarrow \text{P}(\text{OC}_6\text{H}_5)_3 + 3\text{HCl}[/latex]

Similar reactions occur for phosphorus oxychloride, yielding triphenylphosphate:

[latex]\text{OPCl}_3 + 3\text{C}_6\text{H}_5\text{OH} \rightarrow \text{OP}(\text{OC}_6\text{H}_5)_3 + 3\text{HCl}[/latex]

The Phosphate Group

There are several other phosphorus (V) compounds. The most prevalent compounds of phosphorus are derivatives of phosphate (PO43-), a tetrahedral anion. Phosphate is the conjugate base of phosphoric acid, which is produced on a massive scale for use in fertilizers. Since it is triprotic, phosphoric acid converts stepwise to three conjugate bases:

[latex]\text{H}_3\text{PO}_4 + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{H}_2\text{PO}_4^-[/latex] (Ka1 = 7.25 x 10-3)

[latex]\text{H}_2\text{PO}_4^- + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{HPO}_4^{2-}[/latex] (Ka2 = 6.31 x 10-8)

[latex]\text{HPO}_4^{2-} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{PO}_4^{3-}[/latex] (Ka3 = 3.98 x 10-13)

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Phosphoric Acid: Phosphoric acid contains one P=O double bond and three P-O single bonds terminating in acidic OH groups.