Combustion

In complete combustion, the reactant burns in oxygen and produces a limited number of products. When a hydrocarbon burns in oxygen, the reaction will primarily yield carbon dioxide and water. When elements are burned, the products are primarily the most common oxides. Carbon will yield carbon dioxide, sulfur will yield sulfur dioxide, and iron will yield iron(III) oxide. Nitrogen is not considered to be a combustible substance when oxygen is the oxidant. Still, small amounts of various nitrogen oxides (commonly designated NO
_x species) form when the air is the oxidative.

Combustion is not necessarily favorable to the maximum degree of oxidation, and it can be temperature-dependent. For example, sulfur trioxide is not produced quantitatively by the combustion of sulfur. NO_x species appear in significant amounts above about 2,800 °F (1,540 °C), and more is produced at higher temperatures. The amount of NO_x is also a function of oxygen excess.[2]

In most industrial applications and in fires, air is the source of oxygen (O
₂). In the air, each mole of oxygen is mixed with approximately 3.71 mol of nitrogen. Nitrogen does not take part in combustion, but at high temperatures, some nitrogen will be converted to NO
_x (mostly NO, with much smaller amounts of NO
₂). On the other hand, when there is insufficient oxygen to combust the fuel completely, some fuel carbon is converted to carbon monoxide, and some of the hydrogens remain unreacted. A complete set of equations for the combustion of a hydrocarbon in the air, therefore, requires an additional calculation for the distribution of oxygen between the carbon and hydrogen in the fuel.

The amount of air required for complete combustion is known as the "theoretical air" or "stoichiometric air".[3] The amount of air above this value actually needed for optimal combustion is known as the "excess air", and can vary from 5% for a natural gas boiler, to 40% for anthracite coal, to 300% for a gas turbine.[4]

Incomplete combustion will occur when there is not enough oxygen to allow the fuel to react completely to produce carbon dioxide and water. It also happens when the combustion is quenched by a heat sink, such as a solid surface or flame trap. As is the case with complete combustion, water is produced by incomplete combustion; however, carbon and carbon monoxide are produced instead of carbon dioxide.

For most fuels, such as diesel oil, coal, or wood, pyrolysis occurs before combustion. In incomplete combustion, products of pyrolysis remain unburnt and contaminate the smoke with noxious particulate matter and gases. Partially oxidized compounds are also a concern; partial oxidation of ethanol can produce harmful acetaldehyde, and carbon can produce toxic carbon monoxide.

The designs of combustion devices can improve the quality of combustion, such as burners and internal combustion engines. Further improvements are achievable by catalytic after-burning devices (such as catalytic converters) or by the simple partial return of the exhaust gases into the combustion process. Such devices are required by environmental legislation for cars in most countries. They may be necessary to enable large combustion devices, such as thermal power stations, to reach legal emission standards.

The degree of combustion can be measured and analyzed with test equipment. HVAC contractors, firefighters and engineers use combustion analyzers to test the efficiency of a burner during the combustion process. Also, the efficiency of an internal combustion engine can be measured in this way, and some U.S. states and local municipalities use combustion analysis to define and rate the efficiency of vehicles on the road today.

Carbon monoxide is one of the products from incomplete combustion.[5] The formation of carbon monoxide produces less heat than formation of carbon dioxide so complete combustion is greatly preferred especially as carbon monoxide is a poisonous gas. When breathed, carbon monoxide takes the place of oxygen and combines with some of the hemoglobin in the blood, rendering it unable to transport oxygen.[6]

These oxides combine with water and oxygen in the atmosphere, creating nitric acid and sulfuric acids, which return to Earth's surface as acid deposition, or "acid rain." Acid deposition harms aquatic organisms and kills trees. Due to its formation of certain nutrients that are less available to plants such as calcium and phosphorus, it reduces the productivity of the ecosystem and farms. An additional problem associated with nitrogen oxides is that they, along with hydrocarbon pollutants, contribute to the formation of ground level ozone, a major component of smog.[7]

Breathing carbon monoxide causes headache, dizziness, vomiting, and nausea. If carbon monoxide levels are high enough, humans become unconscious or die. Exposure to moderate and high levels of carbon monoxide over long periods is positively correlated with the risk of heart disease. People who survive severe carbon monoxide poisoning may suffer long-term health problems.[8] Carbon monoxide from the air is absorbed in the lungs which then binds with hemoglobin in human's red blood cells. This reduces the capacity of red blood cells that carry oxygen throughout the body.

Smoldering is the slow, low-temperature, flameless form of combustion, sustained by the heat evolved when oxygen directly attacks the surface of a condensed-phase fuel. It is a typically incomplete combustion reaction. Solid materials that can sustain a smoldering reaction include coal, cellulose, wood, cotton, tobacco, peat, duff, humus, synthetic foams, charring polymers (including polyurethane foam) and dust. Common examples of smoldering phenomena are the initiation of residential fires on upholstered furniture by weak heat sources (e.g., a cigarette, a short-circuited wire) and the persistent combustion of biomass behind the flaming fronts of wildfires.

Spontaneous combustion is a type of combustion that occurs by self-heating (increase in temperature due to exothermic internal reactions), followed by thermal runaway (self-heating which rapidly accelerates to high temperatures) and finally, ignition. For example, phosphorus self-ignites at room temperature without the application of heat. Organic materials undergoing bacterial composting can generate enough heat to reach the point of combustion.[9]

Combustion resulting in a turbulent flame is the most used for industrial applications (e.g. gas turbines, gasoline engines, etc.) because the turbulence helps the mixing process between the fuel and oxidizer.

The term 'micro' gravity refers to a gravitational state that is 'low' (i.e., 'micro' in the sense of 'small' and not necessarily a millionth of Earth's normal gravity) such that the influence of buoyancy on physical processes may be considered small relative to other flow processes that would be present at normal gravity. In such an environment, the thermal and flow transport dynamics can behave quite differently than in normal gravity conditions (e.g., a candle's flame takes the shape of a sphere.[10]). Microgravity combustion research contributes to the understanding of a wide variety of aspects that are relevant to both the environment of a spacecraft (e.g., fire dynamics relevant to crew safety on the International Space Station) and terrestrial (Earth-based) conditions (e.g., droplet combustion dynamics to assist developing new fuel blends for improved combustion, materials fabrication processes, thermal management of electronic systems, multiphase flow boiling dynamics, and many others).

Combustion processes that happen in very small volumes are considered micro-combustion. The high surface-to-volume ratio increases specific heat loss. Quenching distance plays a vital role in stabilizing the flame in such combustion chambers.

Generally, the chemical equation for stoichiometric combustion of a hydrocarbon in oxygen is:

${\displaystyle {\ce {C}}_{x}{\ce {H}}_{y}+\left(x+{y \over 4}\right){\ce {O2->}}x{\ce {CO2}}+{y \over 2}{\ce {H2O}}}$

For example, the stoichiometric combustion of methane in oxygen is:

${\displaystyle {\ce {{\underset {methane}{CH4}}+ 2O2 -> CO2 + 2H2O}}}$

If the stoichiometric combustion takes place using air as the oxygen source, the nitrogen present in the air (Atmosphere of Earth) can be added to the equation (although it does not react) to show the stoichiometric composition of the fuel in air and the composition of the resultant flue gas. Treating all non-oxygen components in air as nitrogen gives a 'nitrogen' to oxygen ratio of 3.77, i.e. (100% - O2%) / O2% where O2% is 20.95% vol:

${\displaystyle {\ce {C}}_{x}{\ce {H}}_{y}+z{\ce {O2}}+3.77z{\ce {N2 ->}}x{\ce {CO2}}+{y \over 2}{\ce {H2O}}+3.77z{\ce {N2}}}$

where ${\displaystyle z=x+{y \over 4}}$.

For example, the stoichiometric combustion of methane in air is:

${\displaystyle {\ce {{\underset {methane}{CH4}}+ 2O2}}+7.54{\ce {N2-> CO2 + 2H2O}}+7.54{\ce {N2}}}$

The stoichiometric composition of methane in air is 1 / (1 + 2 + 7.54) = 9.49% vol.

The stoichiometric combustion reaction for C_αH_βO_γ in air:

${\displaystyle {\ce {C_{\mathit {\alpha }}H_{\mathit {\beta }}O_{\mathit {\gamma }}}}+\left(\alpha +{\frac {\beta }{4}}-{\frac {\gamma }{2}}\right)\left({\ce {O_{2}}}+3.77{\ce {N_{2}}}\right)\longrightarrow \alpha {\ce {CO_{2}}}+{\frac {\beta }{2}}{\ce {H_{2}O}}+3.77\left(\alpha +{\frac {\beta }{4}}-{\frac {\gamma }{2}}\right){\ce {N_{2}}}}$

The stoichiometric combustion reaction for C_αH_βO_γS_δ:

${\displaystyle {\ce {C_{\mathit {\alpha }}H_{\mathit {\beta }}O_{\mathit {\gamma }}S_{\mathit {\delta }}}}+\left(\alpha +{\frac {\beta }{4}}-{\frac {\gamma }{2}}+\delta \right)\left({\ce {O_{2}}}+3.77{\ce {N_{2}}}\right)\longrightarrow \alpha {\ce {CO_{2}}}+{\frac {\beta }{2}}{\ce {H_{2}O}}+\delta {\ce {SO_{2}}}+3.77\left(\alpha +{\frac {\beta }{4}}-{\frac {\gamma }{2}}+\delta \right){\ce {N_{2}}}}$

The stoichiometric combustion reaction for C_αH_βO_γN_δS_ε:

${\displaystyle {\ce {C_{\mathit {\alpha }}H_{\mathit {\beta }}O_{\mathit {\gamma }}N_{\mathit {\delta }}S_{\mathit {\epsilon }}}}+\left(\alpha +{\frac {\beta }{4}}-{\frac {\gamma }{2}}+\epsilon \right)\left({\ce {O_{2}}}+3.77{\ce {N_{2}}}\right)\longrightarrow \alpha {\ce {CO_{2}}}+{\frac {\beta }{2}}{\ce {H_{2}O}}+\epsilon {\ce {SO_{2}}}+\left(3.77\left(\alpha +{\frac {\beta }{4}}-{\frac {\gamma }{2}}+\epsilon \right)+{\frac {\delta }{2}}\right){\ce {N_{2}}}}$

The stoichiometric combustion reaction for C_αH_βO_γF_δ:

${\displaystyle {\ce {C_{\mathit {\alpha }}H_{\mathit {\beta }}O_{\mathit {\gamma }}F_{\mathit {\delta }}}}+\left(\alpha +{\frac {\beta -\delta }{4}}-{\frac {\gamma }{2}}\right)\left({\ce {O_{2}}}+3.77{\ce {N_{2}}}\right)\longrightarrow \alpha {\ce {CO_{2}}}+{\frac {\beta -\delta }{2}}{\ce {H_{2}O}}+\delta {\ce {HF}}+3.77\left(\alpha +{\frac {\beta -\delta }{4}}-{\frac {\gamma }{2}}\right){\ce {N_{2}}}}$

Various other substances begin to appear in significant amounts in combustion products when the flame temperature is above about 1600 K. When excess air is used, nitrogen may oxidize to NO and, to a much lesser extent, to NO
₂. CO forms by disproportionation of CO₂, and H
₂ and OH form by disproportionation of H₂O.

For example, when 1 mol of propane is burned with 28.6 mol of air (120% of the stoichiometric amount), the combustion products contain 3.3% O
₂. At 1400 K, the equilibrium combustion products contain 0.03% NO and 0.002% OH. At 1800 K, the combustion products contain 0.17% NO, 0.05% OH, 0.01% CO, and 0.004% H
₂.[11]

Diesel engines are run with an excess of oxygen to combust small particles that tend to form with only a stoichiometric amount of oxygen, necessarily producing nitrogen oxide emissions. Both the United States and European Union enforce limits to vehicle nitrogen oxide emissions, which necessitate the use of special catalytic converters or treatment of the exhaust with urea (see Diesel exhaust fluid).

The incomplete (partial) combustion of a hydrocarbon with oxygen produces a gas mixture containing mainly CO
₂, CO, H₂O, and H
₂. Such gas mixtures are commonly prepared for use as protective atmospheres for the heat-treatment of metals and for gas carburizing.[12] The general reaction equation for incomplete combustion of one mole of a hydrocarbon in oxygen is:

${\displaystyle {\ce {{\underset {fuel}{C_{\mathit {x}}H_{\mathit {y}}}}+{\underset {oxygen}{{\mathit {z}}O2}}->{\underset {carbon\ dioxide}{{\mathit {a}}CO2}}+{\underset {carbon\ monoxide}{{\mathit {b}}CO}}+{\underset {water}{{\mathit {c}}H2O}}+{\underset {hydrogen}{{\mathit {d}}H2}}}}}$

When z falls below roughly 50% of the stoichiometric value, CH
₄ can become an important combustion product; when z falls below roughly 35% of the stoichiometric value, elemental carbon may become stable.

The products of incomplete combustion can be calculated with the aid of a material balance, together with the assumption that the combustion products reach equilibrium.[13][14] For example, in the combustion of one mole of propane (C
₃H
₈) with four moles of O
₂, seven moles of combustion gas are formed, and z is 80% of the stoichiometric value. The three elemental balance equations are:

• Carbon: ${\displaystyle a+b=3}$
• Hydrogen: ${\displaystyle 2c+2d=8}$
• Oxygen: ${\displaystyle 2a+b+c=8}$

These three equations are insufficient in themselves to calculate the combustion gas composition. However, at the equilibrium position, the water-gas shift reaction gives another equation:

${\displaystyle {\ce {CO + H2O -> CO2 + H2}}}$; ${\displaystyle K_{eq}={\frac {a\times d}{b\times c}}}$

For example, at 1200 K the value of K_eq is 0.728.[15] Solving, the combustion gas consists of 42.4% H₂O, 29.0% CO₂, 14.7% H
₂, and 13.9% CO. Carbon becomes a stable phase at 1200 K and 1 atm pressure when z is less than 30% of the stoichiometric value, at which point the combustion products contain more than 98% H
₂ and CO and about 0.5% CH
₄.

Substances or materials which undergo combustion are called fuels. The most common examples are natural gas, propane, kerosene, diesel, petrol, charcoal, coal, wood, etc.

Combustion of a liquid fuel in an oxidizing atmosphere actually happens in the gas phase. It is the vapor that burns, not the liquid. Therefore, a liquid will normally catch fire only above a certain temperature: its flash point. The flash point of liquid fuel is the lowest temperature at which it can form an ignitable mix with air. It is the minimum temperature at which there is enough evaporated fuel in the air to start combustion.

Combustion of gaseous fuels may occur through one of four distinctive types of burning: diffusion flame, premixed flame, autoignitive reaction front, or as a detonation.[16] The type of burning that actually occurs depends on the degree to which the fuel and oxidizer are mixed prior to heating: for example, a diffusion flame is formed if the fuel and oxidizer are separated initially, whereas a premixed flame is formed otherwise. Similarly, the type of burning also depends on the pressure: a detonation, for example, is an autoignitive reaction front coupled to a strong shock wave giving it its characteristic high-pressure peak and high detonation velocity.[16]

The act of combustion consists of three relatively distinct but overlapping phases:

• Preheating phase, when the unburned fuel is heated up to its flash point and then fire point. Flammable gases start being evolved in a process similar to dry distillation.
• Distillation phase or gaseous phase, when the mix of evolved flammable gases with oxygen is ignited. Energy is produced in the form of heat and light. Flames are often visible. Heat transfer from the combustion to the solid maintains the evolution of flammable vapours.
• Charcoal phase or solid phase, when the output of flammable gases from the material is too low for the persistent presence of flame and the charred fuel does not burn rapidly and just glows and later only smoulders.

The kinetic modelling may be explored for insight into the reaction mechanisms of thermal decomposition in the combustion of different materials by using for instance Thermogravimetric analysis.[49]