Chemical Reactions

by Anthony Carpi, Ph.D.

The reaction of two or more elements together results in the formation of a chemical bond between atoms and the formation of a chemical compound (see our Chemical Bonding module). But why do chemicals react together? The reason has to do with the participating atoms' electron configurations (see our The Periodic Table of Elements module).
In the late 1890s, the Scottish chemist Sir William Ramsay discovered the elements helium, neon, argon, krypton, and xenon. These elements, along with radon, were placed in group VIIIA of the periodic table and nicknamed inert (or noble) gases because of their tendency not to react with other elements (see our Periodic Table page). The tendency of the noble gases to not react with other elements has to do with their electron configurations. All of the noble gases have full valence shells; this configuration is a stable configuration and one that other elements try to achieve by reacting together. In other words, the reason atoms react with each other is to reach a state in which their valence shell is filled.

A chemical reaction whenever bonds are formed or broken between molecules. Why certain atoms combine with which other atoms is a complex question which is explained exhaustively by quantum chemistry. This means that the atoms that were joined together in the original substances
break apart and rearrange themselves to make a new substance, one or more substances may be transformed into one or more new
substances. This new substance is quite different from the original substances.
Reactant + Reactant ----> Product
Some products of chemical reactions are heat, light, sound and changes in color.
Chemical reactions are classified according to the kind of change that takes place.
Examples of Chemical Reactions:
    A sparkler contains magnesium. This, when lit, reacts with oxygen in the air and produces light and heat. 
The chemicals inside a rocket, when lit, react with the oxygen in the air and produce heat, light and sound.

All chemical reactions can be placed into one of six categories.  Here they are, in no particular order:

1) Combustion: A combustion reaction is when oxygen combines with another compound to form water and carbon dioxide. These reactions are exothermic, meaning they produce heat. An example of this kind of reaction is the burning of napthalene:

C₁₀H₈ + 12 O₂ ---> 10 CO₂ + 4 H₂O

---

2) Synthesis: A synthesis reaction is when two or more simple compounds combine to form a more complicated one. These reactions come in the general form of:

A + B ---> AB

One example of a synthesis reaction is the combination of iron and sulfur to form iron (II) sulfide:

8 Fe + S₈ ---> 8 FeS

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3) Decomposition: A decomposition reaction is the opposite of a synthesis reaction - a complex molecule breaks down to make simpler ones. These reactions come in the general form:

AB ---> A + B

One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen gas:

2 H₂O ---> 2 H₂ + O₂

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4) Single displacement: This is when one element trades places with another element in a compound. These reactions come in the general form of:

A + BC ---> AC + B

One example of a single displacement reaction is when magnesium replaces hydrogen in water to make magnesium hydroxide and hydrogen gas:

Mg + 2 H₂O ---> Mg(OH)₂ + H₂

---

5) Double displacement: This is when the anions and cations of two different molecules switch places, forming two entirely different compounds. These reactions are in the general form:

AB + CD ---> AD + CB

One example of a double displacement reaction is the reaction of lead (II) nitrate with potassium iodide to form lead (II) iodide and potassium nitrate:

Pb(NO₃)₂ + 2 KI ---> PbI₂ + 2 KNO₃

---

6) Acid-base: This is a special kind of double displacement reaction that takes place when an acid and base react with each other. The H⁺ ion in the acid reacts with the OH^- ion in the base, causing the formation of water. Generally, the product of this reaction is some ionic salt and water:

HA + BOH ---> H₂O + BA

One example of an acid-base reaction is the reaction of hydrobromic acid (HBr) with sodium hydroxide: 

HBr + NaOH ---> NaBr + H₂O

Redox (portmanteau for reduction-oxidation) reactions describe all chemical reactions in which atoms have their oxidation number changed. This can be either a simple redox process, such as the oxidation of carbon to yield carbon dioxide (CO₂) or the reduction of carbon by hydrogen to yield methane (CH₄), or a complex process such as the oxidation of glucose (C₆H₁₂O₆) in the human body through a series of complex electron transfer processes.
Redox reactions, or oxidation-reduction reactions, have a number of similarities to acid-base reactions. Fundamentally, redox reactions are a family of reactions that are concerned with the transfer of electrons between species. Like acid-base reactions, redox reactions are a matched set, that is there cannot be an oxidation reaction without a reduction reaction happening simultaneously. Oxidation refers to the loss of electrons, while reduction refers to the gain of electrons. Each reaction by itself is called a "half-reaction", simply because there must be two half-reactions to form a whole reaction. Thus, in notating redox reactions, chemists typically write out the electrons explicitly:
The term comes from the two concepts of reduction and oxidation. It can be explained in simple terms:

• Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.
• Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.

Though sufficient for many purposes, these descriptions are not precisely correct. Oxidation and reduction properly refer to a change in oxidation number — the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation number, and reduction as a decrease in oxidation number. In practice, the transfer of electrons will always cause a change in oxidation number, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving covalent bonds).
Non-redox reactions, which do not involve changes in formal charge, are known as metathesis reactions.

Acid–base reaction

An acid–base reaction is a chemical reaction that occurs between an acid and a base. Several concepts that provide alternative definitions for the reaction mechanisms involved and their application in solving related problems exist. Despite several differences in definitions, their importance becomes apparent as different methods of analysis when applied to acid–base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent. The first of these scientific concepts of acids and bases was provided by the French chemist Antoine Lavoisier, circa 1776.^[1]

Historic acid–base theories

Lavoisier's oxygen theory of acids

The first scientific concept of acids and bases was provided by Antoine Lavoisier circa 1776. Since Lavoisier's knowledge of strong acids was mainly restricted to oxoacids, such as HNO₃ (nitric acid) and H₂SO₄ (sulphuric acid), which tend to contain central atoms in high oxidation states surrounded by oxygen, and since he was not aware of the true composition of the hydrohalic acids (HF, HCl, HBr, and HI), he defined acids in terms of their containing oxygen, which in fact he named from Greek words meaning "acid-former" (from the Greek οξυς (oxys) meaning "acid" or "sharp" and γεινομαι (geinomai) meaning "engender"). The Lavoisier definition was held as absolute truth for over 30 years, until the 1810 article and subsequent lectures by Sir Humphry Davy in which he proved the lack of oxygen in H₂S, H₂Te, and the hydrohalic acids. However, Davy failed to develop a new theory, concluding that "acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various substances".^[2] One notable modification of oxygen theory was provided by Berzelius, who stated that acids are oxides of nonmetals while bases are oxides of metals.

[edit] Liebig's hydrogen theory of acids

This definition was proposed by Justus von Liebig circa 1838,^[3] based on his extensive works on the chemical composition of organic acids. This finished the doctrinal shift from oxygen-based acids to hydrogen-based acids, started by Davy. According to Liebig, an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal.^[4] Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.^[5]

[edit] Common acid–base theories

[edit] Arrhenius definition

Svante Arrhenius

The Arrhenius definition of acid–base reactions is a development of the hydrogen theory of acids, devised by Svante Arrhenius, which was used to provide a modern definition of acids and bases that followed from his work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution in 1884, and led to Arrhenius receiving the Nobel Prize in Chemistry in 1903 for "recognition of the extraordinary services... rendered to the advancement of chemistry by his electrolytic theory of dissociation".^[6]
As defined by Arrhenius, acid–base reactions are characterized by Arrhenius acids, which dissociate in aqueous solution to form hydrogen ions (H⁺),^[6] and Arrhenius bases, which form hydroxide (OH⁻) ions. More recent IUPAC recommendations now suggest the newer term "hydronium"^[7] be used in favor of the older accepted term "oxonium"^[8] to illustrate reaction mechanisms such as those defined in the Brønsted–Lowry and solvent system definitions more clearly, with the Arrhenius definition serving as a simple general outline of acid–base character.^[6] The Arrhenius definition can be summarised as "Arrhenius acids form hydrogen ions in aqueous solution with Arrhenius bases forming hydroxide ions."
The universal aqueous acid–base definition of the Arrhenius concept is described as the formation of water from hydrogen and hydroxide ions, or hydrogen ions and hydroxide ions from the dissociation of an acid and base in aqueous solution:

H⁺ (aq) + OH⁻ (aq) H₂O

(In modern times, the use of H⁺ is regarded as a shorthand for H₃O⁺, since it is now known that the bare proton H⁺ does not exist as a free species in solution.)^{[citation needed]}
This leads to the definition that in Arrhenius acid–base reactions, a salt and water is formed from the reaction between an acid and a base.^[6] In other words, this is a neutralization reaction.

acid⁺ + base⁻ → salt + water

The positive ion from a base forms a salt with the negative ion from an acid. For example, two moles of the base sodium hydroxide (NaOH) can combine with one mole of sulfuric acid (H₂SO₄) to form two moles of water and one mole of sodium sulfate.

2 NaOH + H₂SO₄ → 2 H₂O + Na₂SO₄

The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions, and refer to the concentration of the solvent ions. Under this definition, pure H₂SO₄ or HCl dissolved in toluene are not acidic, and molten KOH and solutions of sodium amide in liquid ammonia are not alkaline.

[edit] Solvent system definition

One of the limitations of Arrhenius definition was its reliance on water solutions. Edward C. Franklin studied the acid–base reactions in liquid ammonia in 1905 and pointed out the similarities to water-based Arrhenius theory, and Albert F. O. Germann, working with liquid COCl₂, generalized Arrhenius definition to cover aprotic solvents and formulated the solvent system theory in 1925.^[9]
Germann pointed out that in many solvents there is a certain concentration of a positive species, solvonium (earlier lyonium) cations and negative species, solvate (earlier lyate) anions, in equilibrium with the neutral solvent molecules. For example, water and ammonia undergo such dissociation into hydronium and hydroxide, and ammonium and amide, respectively:

2 H₂O H₃O⁺ + OH⁻

2 NH₃ NH+ 4 + NH− 2

Some aprotic systems also undergo such dissociation, such as dinitrogen tetroxide into nitrosonium and nitrate, antimony trichloride into dichloroantimonium and tetrachloroantimonate, and phosgene into chlorocarboxonium and chloride.

N₂O₄ NO⁺ + NO− 3

2 SbCl₃ SbCl+ 2 + SbCl− 4

COCl₂ COCl⁺ + Cl⁻

A solute causing an increase in the concentration of the solvonium ions and a decrease in the solvate ions is defined as an acid and one causing the reverse is defined as a base. Thus, in liquid ammonia, KNH₂ (supplying NH−
2) is a strong base, and NH₄NO₃ (supplying NH+
4) is a strong acid. In liquid sulfur dioxide (SO₂), thionyl compounds (supplying SO²⁺) behave as acids, and sulfites (supplying SO2−
3) behave as bases.
The non-aqueous acid–base reactions in liquid ammonia are similar to the reactions in water:

2 NaNH₂ (base) + Zn(NH₂)₂ (amphiphilic amide) → Na₂[Zn(NH₂)₄]

2 NH₄I (acid) + Zn(NH₂)₂ (amphiphilic amide) → [Zn(NH₃)₄)]I₂

Nitric acid can be a base in liquid sulfuric acid:

HNO₃ (base) + 2 H₂SO₄ → NO+ 2 + H₃O⁺ + 2 HSO− 4

The unique strength of this definition shows in describing the reactions in aprotic solvents, for example in liquid N₂O₄:

AgNO₃ (base) + NOCl (acid) → N₂O₄ (solvent) + AgCl (salt)

Since solvent-system definition depends on the solvent as well as on the compound itself, the same compound can change its role depending on the choice of the solvent. Thus, HClO₄ is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid. This was seen as both a strength and a weakness, since some substances, such as SO₃ and NH₃ were felt to be acidic or basic on their own right. On the other hand, solvent system theory was criticized as too general to be useful; it was felt that there was something intrinsically acidic about hydrogen compounds, not shared by non-hydrogenic solvonium salts.^[2]

[edit] Brønsted–Lowry definition

Main article: Brønsted–Lowry acid–base theory

The Brønsted–Lowry definition, formulated in 1923, independently by Johannes Nicolaus Brønsted in Denmark and Martin Lowry in England, is based upon the idea of protonation of bases through the de-protonation of acids – that is, the ability of acids to "donate" hydrogen ions (H⁺) or protons to bases, which "accept" them.^[10] Unlike the previous definitions, the Brønsted–Lowry definition does not refer to the formation of salt and solvent, but instead to the formation of conjugate acids and conjugate bases, produced by the transfer of a proton from the acid to the base.^[6][10] In this approach, acids and bases are fundamentally different in behavior from salts, which are seen as electrolytes, subject to the theories of Debye, Onsager, and others. An acid and a base react not to produce a salt and a solvent, but to form a new acid and a new base. The concept of neutralization is thus absent.^[2]
According to Brønsted–Lowry definition, an acid is a compound that can donate a proton, and a base is a compound that can receive a proton. An acid–base reaction is, thus, the removal of a hydrogen ion from the acid and its addition to the base.^[11] This does not refer to the removal of a proton from the nucleus of an atom, which would require levels of energy not attainable through the simple dissociation of acids, but to removal of a hydrogen ion (H⁺).
The removal of a proton (hydrogen ion) from an acid produces its conjugate base, which is the acid with a hydrogen ion removed, and the reception of a proton by a base produces its conjugate acid, which is the base with a hydrogen ion added.
For example, the removal of H⁺ from hydrochloric acid (HCl) produces the chloride ion (Cl⁻), the conjugate base of the acid:

HCl → H⁺ + Cl⁻

The addition of H⁺ to the hydroxide ion (OH⁻), a base, produces water (H₂O), its conjugate acid:

H⁺ + OH⁻ → H₂O

Although Brønsted–Lowry acid–base behavior is formally independent of any solvent, it encompasses Arrhenius and solvent system definitions in an unenforced way. For example, protonation of ammonia, a base, gives ammonium ion, its conjugate acid:

H⁺ + NH₃ → NH+ 4

The reaction of ammonia, a base, with acetic acid in absence of water can be described to give ammonium cation, an acid, and acetate anion, a base:

CH₃COOH + NH₃ → NH+ 4 + CH₃COO⁻

This definition also explains the dissociation of water into low concentrations of hydronium and hydroxide ions:

H₂O + H₂O H₃O⁺ + OH⁻

Water, being amphoteric, can act as both an acid and a base; here, one molecule of water acts as an acid, donating a H⁺ ion and forming the conjugate base, OH⁻, and a second molecule of water acts as a base, accepting the H⁺ ion and forming the conjugate acid, H₃O⁺.
Acid dissociation and acid hydrolysis are seen to be entirely similar phenomena:

HCl (acid) + H₂O (base) H₃O⁺ (acid) + Cl⁻ (base)

NH+ 4 (acid) + H₂O (base) H₃O⁺ (acid) + NH₃ (base)

as are basic dissociation and basic hydrolysis:

NH₃ (base) + H₂O (acid) NH+ 4 (acid) + OH⁻ (base)

CH₃COO⁻ (base) + H₂O (acid) CH₃COOH (acid) + OH⁻ (base)

Thus, the general formula for acid–base reactions according to the Brønsted–Lowry definition is:

AH + B → BH⁺ + A⁻

where AH represents the acid, B represents the base, BH⁺ represents the conjugate acid of B, and A⁻ represents the conjugate base of AH.
Although Brønsted–Lowry calls hydrogen-containing substances like HCl acids, KOH and KNH₂ are not bases but salts containing the bases OH⁻ and NH−
2. Also, some substances, which many chemists considered to be acids, such as SO₃ or BCl₃, are excluded from this classification due to lack of hydrogen. Gilbert Lewis wrote in 1938, "To restrict the group of acids to those substances which contain hydrogen interferes as seriously with the systematic understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen."^[2]

[edit] Lewis definition

Further information: Lewis acids and bases

The hydrogen requirement of Arrhenius and Brønsted–Lowry was removed by the Lewis definition of acid–base reactions, devised by Gilbert N. Lewis in 1923,^[12] in the same year as Brønsted–Lowry, but it was not elaborated by him until 1938.^[2] Instead of defining acid–base reactions in terms of protons or other bonded substances, the Lewis definition defines a base (referred to as a Lewis base) to be a compound that can donate an electron pair, and an acid (a Lewis acid) to be a compound that can receive this electron pair.^[13]
In this system, an acid does not exchange atoms with a base, but combines with it. For example, consider this classical aqueous acid–base reaction:

HCl (aq) + NaOH (aq) → H₂O (l) + NaCl (aq)

The Lewis definition does not regard this reaction as the formation of salt and water or the transfer of H⁺ from HCl to OH⁻. Instead, it regards the acid to be the H⁺ ion itself, and the base to be the OH⁻ ion, which has an unshared electron pair. Therefore, the acid–base reaction here, according to the Lewis definition, is the donation of the electron pair from OH⁻ to the H⁺ ion. This forms a covalent bond between H⁺ and OH⁻, thus producing water (H₂O).
By treating acid–base reactions in terms of electron pairs instead of specific substances, the Lewis definition can be applied to reactions that do not fall under other definitions of acid–base reactions. For example, a silver cation behaves as an acid with respect to ammonia, which behaves as a base, in the following reaction:

Ag⁺ + 2 :NH₃ → [H₃N:Ag:NH₃]⁺

The result of this reaction is the formation of an ammonia–silver adduct.
In reactions between Lewis acids and bases, there is the formation of an adduct^[13] when the highest occupied molecular orbital (HOMO) of a molecule, such as NH₃ with available lone electron pair(s) donates lone pairs of electrons to the electron-deficient molecule's lowest unoccupied molecular orbital (LUMO) through a co-ordinate covalent bond; in such a reaction, the HOMO-interacting molecule acts as a base, and the LUMO-interacting molecule acts as an acid.^[13] In highly-polar molecules, such as boron trifluoride (BF₃),^[13] the most electronegative element pulls electrons towards its own orbitals, providing a more positive charge on the less-electronegative element and a difference in its electronic structure due to the axial or equatorial orbiting positions of its electrons, causing repulsive effects from lone pair – bonding pair (Lp–Bp) interactions between bonded atoms in excess of those already provided by bonding pair – bonding pair (Bp–Bp) interactions.^[13] Adducts involving metal ions are referred to as co-ordination compounds.^[13]

[edit] Other acid–base theories

[edit] Usanovich definition

Simultaneously with Lewis, a Soviet chemist Mikhail Usanovich from Tashkent, developed a general theory that does not restrict acidity to hydrogen-containing compounds, but his approach, published in 1938, was even more general than Lewis theory.^[2] Usanovich's theory can be summarized as defining an acid as anything that accepts negative species or donates positive ones, and a base as the reverse. This pushed the concept of acid–base reactions to its logical limits, and even redefined the concept of redox (oxidation-reduction) as a special case of acid-base reactions, and so did not become widespread, despite being easier to understand than Lewis theory, which required detailed familiarity with atomic structure. Some examples of Usanovich acid-base reactions include:

Na₂O (base) + SO₃ (acid) → 2 Na⁺ + SO2− 4 (species exchanged: anion O²⁻)

3 (NH₄)₂S (base) + Sb₂S₃ (acid) → 6 NH+ 4 + 2 SbS2− 4 (species exchanged: anion S²⁻)

Na (base) + Cl (acid) → Na⁺ + Cl⁻ (species exchanged: electron)

[edit] Lux–Flood definition

This acid–base theory was a revival of oxygen theory of acids and bases, proposed by German chemist Hermann Lux^[14][15] in 1939, further improved by Håkon Flood circa 1947^[16] and is still used in modern geochemistry and electrochemistry of molten salts. This definition describes an acid as an oxide ion (O²⁻) acceptor and a base as an oxide ion donor. For example:^[17]

MgO (base) + CO₂ (acid) → MgCO₃

CaO (base) + SiO₂ (acid) → CaSiO₃

NO− 3 (base) + S₂O2− 7 (acid) → NO+ 2 + 2 SO2− 4

[edit] Pearson definition

Main article: HSAB theory

In 1963^[18] Ralph Pearson proposed an advanced qualitative concept known as Hard Soft Acid Base principle, later made quantitative with help of Robert Parr in 1984. 'Hard' applies to species that are small, have high charge states, and are weakly polarizable. 'Soft' applies to species that are large, have low charge states and are strongly polarizable. Acids and bases interact, and the most stable interactions are hard–hard and soft–soft. This theory has found use in organic and inorganic chemistry.

[edit] Acid–alkali reaction

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It has been suggested that this article or section be merged into Neutralization (chemistry). (Discuss) Proposed since January 2010.

An acid–alkali reaction is a special case of an acid–base reaction, where the base used is also an alkali. When an acid reacts with an alkali it forms a metal salt and water. Acid–alkali reactions are also a type of neutralization reaction.
In general acid–alkali reactions can be simplified to

OH⁻(aq) + H⁺(aq) → H₂O

by omitting spectator ions.
Acids are generally pure substances which contain hydrogen ions (H⁺) or cause them to be produced in solutions. Hydrochloric acid (HCl) and sulfuric acid (H₂SO₄) are common examples. In water, these break apart into ions:

HCl → H⁺(aq) + Cl⁻(aq)

H₂SO₄ → H⁺(aq) + HSO− 4(aq)

An alkali is a base, more precisely a base which contains a metal from column 1 or 2 of the periodic table (the alkali metals or the alkaline earth metals). Alkalis may be defined as soluble bases, which means they must be able to dissolve in water. Bases generally are defined as substances which contain hydroxide ion (OH⁻) or produce it in solution. Therefore, one may also speak of hydroxide bases which dissolve in water, and thus these would also be alkalis. Some examples, then, of alkalis would be sodium hydroxide (NaOH), potassium hydroxide (KOH), magnesium hydroxide (Mg(OH)₂), and calcium hydroxide (Ca(OH)₂). Note that only hydroxides with an alkali metal—column 1—are very soluble in water; hydroxides with an alkaline earth metal—column 2—are not as soluble. Some sources^[19] will even say the alkaline earth metal hydroxides are insoluble.
To produce hydroxide ions in water, the alkali breaks apart into ions as below:

NaOH → Na⁺(aq) + OH⁻(aq)

However, alkalies may also have a broader definition which includes carbonates (CO2−
3) bonded to a column 1 metal, an ammonium ion (NH+
4), or an amine (NH_x radical) as the positive ion. Examples of alkalis would then also include Li₂CO₃, Na₂CO₃, and (NH₄)₂CO₃.
There seems to be conflicting information on acid-base reactions being neutralization reactions. Some sources define a neutralization reaction as the reaction between an acid and a base which produces a salt and water. Yet in the book Chemical Misconceptions: Prevention, Diagnosis and Cure by K. Tabor (2002), it is noted that “the term neutralization is usually reserved for acid–alkali reactions.” Thus this does not make acid–alkali a type of neutralization reaction, but the only kind of neutralization reaction.
There are many uses of neutralization reactions which are acid-alkali reactions. A very common use is antacid tablets. These are designed to neutralize excess stomach acid (HCl) which may be causing discomfort in the stomach or lower esophagus. Also in the digestive tract, neutralization reactions are used when food is moved from the stomach to the intestines. In order for the nutrients to be absorbed through the intestinal wall, an alkaline environment is needed, so the pancreas produce an antacid bicarbonate to cause this transformation to occur.^[20]
Another common use, though perhaps not as widely known, is in fertilizers and control of soil pH. Slaked lime (calcium hydroxide) or limestone (calcium carbonate) may be worked into soil that is too acidic for plant growth.^[21] Fertilizers which improve plant growth are made by neutralizing sulfuric acid (H₂SO₄) or nitric acid (HNO₃) with ammonia gas (NH₃) making ammonium sulfate or ammonium nitrate. These are salts utilized in the fertilizer.^[22]
Industrially, a by-product of the burning of coal, sulfur dioxide gas may combine with water vapor in the air to eventually produce sulfuric acid, which falls as acid rain. To prevent the sulfur dioxide from being released, a device known as a scrubber gleans the gas from smoke stacks. This device first blows calcium carbonate into the combustion chamber where it decomposes into calcium oxide (lime) and carbon dioxide. This lime then reacts with the sulfur dioxide produced forming calcium sulfite. A suspension of lime is then injected into the mixture to produce a slurry, which removes the calcium sulfite and any remaining unreacted sulfur dioxide.