Mechanism of meat curing – the wonder of nitrogen

Mechanism of meat curing – the wonder of nitrogen
By:  Eben van Tonder
4 September 2016


In these articles we examine the mechanics of meat curing so that we can ensure factory conditions and processing steps that favour curing.

In our previous article we set the historical context of our understanding of curing mechanisms as it relates to colour development in cooked cured meat.  Here we deal with the reactivity of nitrogen and how it changes into a form that we use in meat curing.


This article follows the formation of some of the important nitrogen compounds and molecules starting with nitrogen gas in the atmosphere, the formation of nitric oxide, nitrogen dioxide, nitric acid and ammonia.  We briefly looks at some of the mechanisms that change nitrogen gas into a usable form for meat curing as well as the reactive nature of nitrogen.


Farmers in the 1880’2 were universally well informed about nitrogen and its value in soil.  It is a gas that forms part of our atmosphere.  We estimate, at least 4/5th.  The rest, roughly 20%, is oxygen.  (2)  (Marion Record, 1887, p3: About Nitrogen)

An article from the Marion Record in 1887, reminded us that the role and effect of nitrogen in human and animal tissue is relevant, not just to the living but also to pork from which we make bacon.  The art of bacon is the art of the manipulation of the properties of meat through nitrogen, sodium and chloride.

The curing of meat revolves around nitrogen and it is helpful to know a bit more about the reactivity of this unique chemical element.  It forms the link between fertilizers, food processing and war since the same power fires bullet, provides nutrition to plants and cures meat for future consumption.


Nitrogen gas molecule (N2)

Nitrogen was so named by the early chemists as the generator of nitre.  Nitre is also called saltpeter.

Nitrogen was independently discovered by two scientists.  In 1772, by the Scottish physicist, Daniel Rutherford  (Marion Record, 1887, p3: About Nitrogen) and in the early 1770’s by a Swedish chemist, Carl Scheele.  “Rutherford named his discovery “noxious air,” because animals were not able to breath in it.  Scheele called it “foul air.”  (Farndon, J, 1999: 9)

It was Antoine Lavoisier (1743 – 1794) who realized that air was basically a mixture between two gasses, oxygen and nitrogen.  He burned mercury in a closed jar and found that a 5th of the air combined with the mercury to form a red powder, mercury oxide.  No matter what he did, the rest stayed a gas.  Mice died in it and a candle could not burn in it.  “Lavoisier decided that air is made of two gases.  One, which he called oxygen, was the gas that burned with the mercury.  The other he called azote from the Greek for ‘no life.’  It later came to be known as nitrogen, because it can be generated from niter, the common name for sodium or potassium nitrate or saltpeter” (Farndon, J, 1999: 9)

At ambient temperature, the gas, nitrogen, is an inert molecule.  It is however one of only two elements that can occur in eight oxidation states, the second one being carbon. (Honikel, 2007)

“The outer shell of five electrons (CodeCogsEqn (19)) can take up three additional electrons giving the nitrogen an ‘‘oxidation status’’ CodeCogsEqn (20) as it exists in ammonia (CodeCogsEqn (21)) or amines or it can release five electrons forming CodeCogsEqn (23) as it exists in nitrate CodeCogsEqn (24).” “This is the reason for the variability and wide range of carbon compounds (organic matter) but also for the complexity of nitrogen reactivity. The latter is shown below.”  (Honikel, 2007)  Below one can see the most important nitrogen compounds in their different states of oxidation.

oxidation states of nitrogen.png
Oxidation states of nitrogen (Honikel)

It is enlightening to understand one of the ways that nitrogen change from an inert gas to a form that is used as food for plants and from our vantage point, ends up curing meat.

Two of the ways it enters the food chain is through the power of lightning and the small microorganisms.  Let’s first look at nitrogen that falls from the skies.


Nitric Oxide (NO)

Nitrogen gas exists as two atoms, tightly bound in one molecule (N2).  The bonds between the atoms are so strong that it doesn’t normally react with anything else.  Lightning provides enough energy to break these strong bonds which now makes the nitrogen available to react with other elements.  (Farndon, J, 1999: 10)

One of these elements is oxygen.  When they react, they form nitrogen monoxide (NO).  Nitrogen monoxide is a colourless gas, also called nitric oxide or nitrogen oxide.  The nitric oxide is heated due to the energy from the lightning flash that created it.  (Farndon, J, 1999: 10)

The reaction is written as follows:

N2 (g) + O2 (g)  lightning —> 2NO (g)


nitrogen dioxide
Nitrogen dioxide. (NO2)

Other sources of nitric oxide, besides lightning, are certain bacteria and volcanos.  (Air Quality Guidelines, 2000:  chapter 7).  As it cools down, it reacts further with the oxygen molecules around it to form nitrogen dioxide.  One nitrogen atom attached to two oxygen atoms forms nitrogen dioxide. “It is a poisonous, brown, acidic, pungent gas”.  (Farndon, J, 1999: 12)  Nitrogen dioxide is however mainly formed in the atmosphere through it’s a reaction with ozone (O3).

Like nitrogen, oxygen occurs as two oxygen atoms, bound in one molecule.  Ultra-violet light and lightning cause the two tightly bound oxygen atoms to separate and react, either with other single atom oxygen molecules or with more stable two atom oxygen molecules.  In the latter case, three oxygen atoms are bound into one molecule (O3). (3)  (Wikipedia, Ozone)  It is not very stable and quickly breaks down into one oxygen atom (O) and or two oxygen atom molecules or it reacts with nitric oxide to form nitrogen dioxide.  (Huffman, R. E.; 1992: 210) (Air Quality Guidelines, 2000:  chapter 7)

The reaction occurs as follows:

NO (g) + 1/2O2 (g) —> NO2 (g)


Nitric Acid (HNO3)

Nitrogen Dioxide (NO2) reacts with more oxygen and rain drops to form nitric acid (HNO3) which falls to earth and enters the soil to provide nutrients for plants.  (Ramakrishna, A.; 2014: 14) Nitric acid (HNO3) is also known as aqua fortis and spirit of niter.  (Wikipedia, Nitric Acid)

This puzzling phrase, “spirit of” something seems to have been used generally by chemists when they did not really know what it was.  The particular phrase, “spirit of niter” was puzzling to even  Robert Boyle in the 1650’s and 60’s.  (Rattansi, P.;1994: 66)

The reaction occurs as follows:

3NO2 (g) + H2O —> 2HNO3 (aq) + NO (g)

Nitric acid is highly reactive and combines with salts in the soil, converting it to nitrates which in turn become food for the plants.  (Ramakrishna, A.; 2014: 14)  It is this reaction of nitric acid with salts that create sodium nitrate  or calcium nitrate or potassium nitrate that are used as fertilizer or in gunpowder or to cure bacon.

It has been discovered that curing happens much faster if nitrite is used directly.  Bacteria are responsible for changing nitrate to nitrite when it is injected into meat as a curing agent, just as it is done by bacteria in soil.  Nitrite (NO2) is the same as nitrate (NO3), with one less oxygen atom.  By using nitrite directly, curing is accomplished much faster since the reduction to nitrite takes time.


Ammonia (NH3)

Nitrogen comes into our lives from the atmosphere, but despite the fact that “nitrogen oxides trapped in rocks and sediments probably represent a larger total quantity of nitrogen, this nitrogen, for the most part, is not accessible to living organisms.”  (Igarashi, Y. and Seefeldt, C. L..  2003)   Most nitrogen enters our world through special bacteria that take nitrogen from the atmosphere and combine it with another important chemical element, hydrogen, to produce ammonia.  (

There are many bacteria who achieve this conversion through various means, but a common denominator is that they all use the most interesting enzyme, nitrogenase.  It is this enzyme that is responsible for changing N2 to ammonia.  The general Nreduction reaction catalyzed by these enzymes is typically presented as follows:

N2 + 8 e− + 16 ATP + 8 H+ → 2 NH3 + 16 ADP + 16 Pi + H2  (Igarashi, Y. and Seefeldt, C. L..  2003)

This amazing enzyme has the ability to break apart the very strong N2 molecule and form ammonia.  “Ammonia is easily manipulated by biological cells and by converting it into ammonium (NH4+) and other compounds such as nitrate and nitrites.”   (Dincer, I. and Zamfirescu, C.; 2011: 706)  Interestingly enough, a small amount of ammonia is also produced through pressure and energy from lightning.  (Krasny, M. E.; 2003: 46)

Bacteria with this remarkable ability are found in fresh water, soil and in seawater.  A few of these bacteria live in a special relationship with plants where both benefits in special ways.  The bacteria live in the roots and supply the plant with nitrogen.  In turn, the plant supplies the bacteria with sugars and other carbon compounds.  Examples of these plants are alfalfa, clover, peas, peanuts and beans.  (Krasny, M. E.; 2003: 46)

Ammonium (NH4+) is taken up by the plants and incorporated in amino acids, the building blocks of proteins.  When animals or humans eat the plants, the nitrogen is taken up in their bodies in the form of amino acids and proteins.  (Krasny, M. E.; 2003: 46)

“Similar to Carbon, organic nitrogen is returned to the atmosphere when plants and animals die and are decomposed.  Bacteria first break protein and amino acids back down into ammonium.”  The process now becomes complicated as some ammonium is taken up again by plants and used by the plants to build amino acids and protein.  Some are broken down further by bacteria into nitrite (NO2)and nitrate (NO3).  Nitrate (NO3) itself can be taken up directly by the plants. Some of the nitrate are transformed by bacteria into gaseous nitrous oxide (N2O), nitric oxide (NO), or nitrogen gas (N2), which are released into the air.  Some nitrate makes its way into streams, lakes and groundwater.  (Krasny, M. E.; 2003: 46)


Nitrate, nitrite, nitric acid and nitrous acid take center stage in our chemical sequence development from nitrite to nitric oxide (NO) which now becomes key in the subsequent articles.


Air Quality Guidelines – Second Edition.  2000. Published by the WHO, Regional Office for Europe, Copenhagen, Denmark.

Butcher, S. S.. et al.  1992.  Global biogeochemical cycles. Academic Press, Inc.

Dincer, I. and Zamfirescu, C.  2011.  Sustainable Energy Systems and Applications.  Springer Science + BusinessMedia, LLC.

Farndon, J.  1999.  The Elements, Nitrogen.  Marshall Cavendish Corporation.

Honikel, K-O.  31 May 2007.  The use and control of nitrate and nitrite for the processing of meat products.  Science Direct.  Meat Science 78 (2008) 68–76. Elsevier Ltd.

Huffman, R. E..  1992.  Atmospheric Ultraviolet Remote Sensing.  Academic Press, Inc.

Igarashi, Y. and Seefeldt, C. L..  2003.  Nitrogen Fixation: The Mechanism of the Mo-Dependent Nitrogenase.  Article from Critical Reviews in Biochemistry and Molecular Biology, 38:351–384.  Robert. Taylor and Francis Inc.

Krasny, M. E..  2003.  Invasion Ecology.  NSTA Press.

Langa, S. L..  1999.  The Impact of Nitrogen Deposition on Natural and Semi-Natural Ecosystems.  Springer Science+Business Media.

Marion Record, Marion, Kansas.  Friday, 15 July 1887. About nitrogen, p3

Ramakrishna, A.  2014.  Goyal’s IIT FOUNDATION COURSE CHEMISTRY.  Roshan Lal Goyal for Goyal Brothersn Prakashan.

Rattansi, P..  1994.  Alchemy and Chemistry in the 16th and 17th Centuries.  Kluwer Academic Publishers.  Atmospheric Science. Earth’s atmospheric air.

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