Redox Chemistry

From Wikipedia, the free encyclopedia

Sodium and fluorine bonding ionically to form sodium fluoride. Sodium loses its outer electron to give it a stable electron configuration, and this electron enters the fluorine atom exothermically. The oppositely charged ions are then attracted to each other. The sodium is oxidized, and the fluorine is reduced.

The two parts of a redox reaction

Rusting iron

A bonfire

Redox (portmanteau of reduction and oxidation) reactions include all chemical reactions in which atoms have their oxidation state changed; in general, redox reactions involve the transfer of electrons between species.

The term "redox" comes from two concepts involved with electron transfer: reduction and oxidation.^[1] 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.

Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, these are only specific examples of a more general concept of reactions involving electron transfer.

Redox reactions, or oxidation-reduction reactions, have a number of similarities to acid–base reactions. Like acid–base reactions, redox reactions are a matched set, that is, there cannot be an oxidation reaction without a reduction reaction happening simultaneously. The oxidation alone and the reduction alone are each called a half-reaction, because two half-reactions always occur together to form a whole reaction. When writing half-reactions, the gained or lost electrons are typically included explicitly in order that the half-reaction be balanced with respect to electric charge.

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

There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide (CO
2) or the reduction of carbon by hydrogen to yield methane (CH₄), and more complex processes such as the oxidation of glucose (C₆H₁₂O₆) in the human body through a series of complex electron transfer processes.

Etymology

"Redox" is a portmanteau of "reduction" and "oxidation".

The word oxidation originally implied reaction with oxygen to form an oxide, since dioxygen (O₂ (g)) was historically the first recognized oxidizing agent. Later, the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. Ultimately, the meaning was generalized to include all processes involving loss of electrons.

The word reduction originally referred to the loss in weight upon heating a metallic ore such as a metal oxide to extract the metal. In other words, ore was "reduced" to metal. Antoine Lavoisier (1743-1794) showed that this loss of weight was due to the loss of oxygen as a gas. Later, scientists realized that the metal atom gains electrons in this process. The meaning of reduction then became generalized to include all processes involving gain of electrons. Even though "reduction" seems counter-intuitive when speaking of the gain of electrons, it might help to think of reduction as the loss of oxygen, which was its historical meaning.

The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes respectively when they occur at electrodes.^[2] These words are analogous to protonation and deprotonation, but they have not been widely adopted by chemists.

The term "hydrogenation" could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions, especially in organic chemistry and biochemistry. But, unlike oxidation, which has been generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance (e.g., the hydrogenation of unsaturated fats in saturated fats, R-CH=CH-R + H₂ → R-CH₂-CH₂-R).

Oxidizing and reducing agents

In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, and the oxidant or oxidizing agent gains electrons and is reduced. The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidized form, e.g., Fe²⁺/Fe³⁺.

Oxidizers

Substances that have the ability to oxidize other substances (cause them to lose electrons) are said to be oxidative or oxidizing and are known as oxidizing agents, oxidants, or oxidizers. That is, the oxidant (oxidizing agent) removes electrons from another substance, and is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is also called an electron acceptor, hence the name.  

Oxygen is the quintessential oxidizer.
Oxidants are usually chemical substances with elements in high oxidation states (e.g., H
2O
2, MnO−
4, CrO
3, Cr
2O2−
7, OsO
4), or else highly electronegative elements (O₂, F₂, Cl₂, Br₂) that can gain extra electrons by oxidizing another substance.

Reducers

Substances that have the ability to reduce other substances (cause them to gain electrons) are said to be reductive or reducing and are known as reducing agents, reductants, or reducers. The reductant (reducing agent) transfers electrons to another substance, and is thus itself oxidized. And, because it "donates" electrons, the reducing agent is also called an electron donor. Electron donors can also form charge transfer complexes with electron acceptors.

Reductants in chemistry are very diverse. Electropositive elemental metals, such as lithium, sodium, magnesium, iron, zinc, and aluminium, are good reducing agents. These metals donate or give away electrons readily. Hydride transfer reagents, such as NaBH₄ and LiAlH₄, are widely used in organic chemistry,^[3][4] primarily in the reduction of carbonyl compounds to alcohols. Another method of reduction involves the use of hydrogen gas (H₂) with a palladium, platinum, or nickel catalyst. These catalytic reductions are used primarily in the reduction of carbon-carbon double or triple bonds.

Standard electrode potentials (reduction potentials)

Each half-reaction has a standard electrode potential (E⁰_cell), which is equal to the potential difference (or voltage) (E⁰_cell) at equilibrium under standard conditions of an electrochemical cell in which the cathode reaction is the half-reaction considered, and the anode is a standard hydrogen electrode where hydrogen is oxidized: ½ H₂ → H⁺ + e⁻.

The electrode potential of each half-reaction is also known as its reduction potential E⁰_red, or potential when the half-reaction takes place at a cathode. The reduction potential is a measure of the tendency of the oxidizing agent to be reduced. Its value is zero for H⁺ + e⁻ → ½ H₂ by definition, positive for oxidizing agents stronger than H⁺ (e.g., +2.866 V for F₂) and negative for oxidizing agents that are weaker than H⁺ (e.g., –0.763 V for Zn²⁺).^[5]

For a redox reaction that takes place in a cell, the potential difference E⁰_cell = E⁰_cathode – E⁰_anode

However, the potential of the reaction at the anode was sometimes expressed as an oxidation potential, E⁰_ox = – E⁰. The oxidation potential is a measure of the tendency of the reducing agent to be oxidized, but does not represent the physical potential at an electrode. With this notation, the cell voltage equation is written with a plus sign E⁰_cell = E⁰_cathode + E⁰_{ox (anode)}

Examples of redox reactions

Illustration of a redox reaction

A good example is the reaction between hydrogen and fluorine in which hydrogen is being oxidized and fluorine is being reduced:

H
2 + F
2 → 2 HF

We can write this overall reaction as two half-reactions:
the oxidation reaction:

H
2 → 2 H⁺ + 2 e⁻

and the reduction reaction:

F
2 + 2 e⁻ → 2 F⁻

Analyzing each half-reaction in isolation can often make the overall chemical process clearer.
Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above).

Elements, even in molecular form, always have an oxidation state of zero. In the first half-reaction, hydrogen is oxidized from an oxidation state of zero to an oxidation state of +1. In the second half-reaction, fluorine is reduced from an oxidation state of zero to an oxidation state of −1.
When adding the reactions together the electrons are canceled:

H 2	→	2 H+ + 2 e−
F 2 + 2 e−	→	2 F−
H 2 + F 2	→	2 H+ + 2 F−

And the ions combine to form hydrogen fluoride:

2 H⁺ + 2 F⁻ → 2 HF

The overall reaction is:

H
2 + F
2 → 2 HF

Metal displacement

A redox reaction is the force behind an electrochemical cell like the Galvanic cell pictured. The battery is made out of a zinc electrode in a ZnSO₄ solution connected with a wire and a porous disk to a copper electrode in a CuSO₄ solution.

In this type of reaction, a metal atom in a compound (or in a solution) is replaced by an atom of another metal. For example, copper is deposited when zinc metal is placed in a copper(II) sulfate solution:

Zn(s)+ CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

In the above reaction, zinc metal displaces the copper(II) ion from copper sulfate solution and thus liberates free copper metal.

The ionic equation for this reaction is:

Zn + Cu²⁺ → Zn²⁺ + Cu

As two half-reactions, it is seen that the zinc is oxidized:

Zn → Zn²⁺ + 2 e−

And the copper is reduced:

Cu²⁺ + 2 e− → Cu

Other examples

• The reduction of nitrate to nitrogen in the presence of an acid (denitrification):

2 NO₃⁻ + 10 e⁻ + 12 H⁺ → N₂ + 6 H₂O

• The combustion of hydrocarbons, such as in an internal combustion engine, which produces water, carbon dioxide, some partially oxidized forms such as carbon monoxide, and heat energy. Complete oxidation of materials containing carbon produces carbon dioxide.

• In organic chemistry, the stepwise oxidation of a hydrocarbon by oxygen produces water and, successively, an alcohol, an aldehyde or a ketone, a carboxylic acid, and then a peroxide.

Corrosion and rusting

Oxides, such as iron(III) oxide or rust, which consists of hydrated iron(III) oxides Fe₂O₃·nH₂O and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)₃), form when oxygen combines with other elements

Iron rusting in pyrite cubes

• The term corrosion refers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen. Rusting, the formation of iron oxides, is a well-known example of electrochemical corrosion; it forms as a result of the oxidation of iron metal. Common rust often refers to iron(III) oxide, formed in the following chemical reaction:

4Fe + 3O₂ → 2Fe₂O₃

• The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:

Fe²⁺ → Fe³⁺ + e⁻

H₂O₂ + 2 e⁻ → 2 OH⁻

Overall equation:

2 Fe²⁺ + H₂O₂ + 2 H⁺ → 2 Fe³⁺ + 2 H₂O

Redox reactions in industry

Cathodic protection is a technique used to control the corrosion of a metal surface by making it the cathode of an electrochemical cell. A simple method of protection connects protected metal to a more easily corroded "sacrificial anode" to act as the anode. The sacrificial metal instead of the protected metal, then, corrodes. A common application of cathodic protection is in galvanized steel, in which a sacrificial coating of zinc on steel parts protects them from rust.

The primary process of reducing ore at high temperature to produce metals is known as smelting.

Oxidation is used in a wide variety of industries such as in the production of cleaning products and oxidizing ammonia to produce nitric acid, which is used in most fertilizers.

Redox reactions are the foundation of electrochemical cells.

The process of electroplating uses redox reactions to coat objects with a thin layer of a material, as in chrome-plated automotive parts, silver plating cutlery, and gold-plated jewelry.

The production of compact discs depends on a redox reaction, which coats the disc with a thin layer of metal film.^{[clarification needed]}

Redox reactions in biology

Top: ascorbic acid (reduced form of Vitamin C)
Bottom: dehydroascorbic acid (oxidized form of Vitamin C)

Many important biological processes involve redox reactions.

Cellular respiration, for instance, is the oxidation of glucose (C₆H₁₂O₆) to CO₂ and the reduction of oxygen to water. The summary equation for cell respiration is:

C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O

The process of cell respiration also depends heavily on the reduction of NAD⁺ to NADH and the reverse reaction (the oxidation of NADH to NAD⁺). Photosynthesis and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cell respiration:

6 CO₂ + 6 H₂O + light energy → C₆H₁₂O₆ + 6 O₂

Biological energy is frequently stored and released by means of redox reactions. Photosynthesis involves the reduction of carbon dioxide into sugars and the oxidation of water into molecular oxygen. The reverse reaction, respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce nicotinamide adenine dinucleotide (NAD⁺), which then contributes to the creation of a proton gradient, which drives the synthesis of adenosine triphosphate (ATP) and is maintained by the reduction of oxygen. In animal cells, mitochondria perform similar functions. See the Membrane potential article.

Free radical reactions are redox reactions that occur as a part of homeostasis and killing microorganisms, where an electron detaches from a molecule and then reattaches almost instantaneously. Free radicals are a part of redox molecules and can become harmful to the human body if they do not reattach to the redox molecule or an antioxidant. Unsatisfied free radicals can spur the mutation of cells they encounter and are, thus, causes of cancer.

The term redox state is often used to describe the balance of GSH/GSSG, NAD⁺/NADH and NADP⁺/NADPH in a biological system such as a cell or organ. The redox state is reflected in the balance of several sets of metabolites (e.g., lactate and pyruvate, beta-hydroxybutyrate, and acetoacetate), whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and sepsis. Redox mechanism also control some cellular processes. Redox proteins and their genes must be co-located for redox regulation according to the CoRR hypothesis for the function of DNA in mitochondria and chloroplasts.

Redox cycling

A wide variety of aromatic compounds are enzymatically reduced to form free radicals that contain one more electron than their parent compounds. In general, the electron donor is any of a wide variety of flavoenzymes and their coenzymes. Once formed, these anion free radicals reduce molecular oxygen to superoxide, and regenerate the unchanged parent compound. The net reaction is the oxidation of the flavoenzyme's coenzymes and the reduction of molecular oxygen to form superoxide. This catalytic behavior has been described as futile cycle or redox cycling.

Examples of redox cycling-inducing molecules are the herbicide paraquat and other viologens and quinones such as menadione.^[6]

Redox reactions in geology

Mi Vida uranium mine, near Moab, Utah. The alternating red and white/green bands of sandstone correspond to oxidized and reduced conditions in groundwater redox chemistry.

In geology, redox is important to both the formation of minerals and the mobilization of minerals, and is also important in some depositional environments. In general, the redox state of most rocks can be seen in the color of the rock. The rock forms in oxidizing conditions, giving it a red color. It is then "bleached" to a green—or sometimes white—form when a reducing fluid passes through the rock. The reduced fluid can also carry uranium-bearing minerals. Famous examples of redox conditions affecting geological processes include uranium deposits and Moqui marbles.

Balancing redox reactions

Describing the overall electrochemical reaction for a redox process requires a balancing of the component half-reactions for oxidation and reduction. In general, for reactions in aqueous solution, this involves adding H⁺, OH⁻, H₂O, and electrons to compensate for the oxidation changes.

Acidic media

In acidic media, H+ ions and water are added to half-reactions to balance the overall reaction.
For instance, when manganese(II) reacts with sodium bismuthate:

Unbalanced reaction:	Mn2+(aq) + NaBiO 3(s) → Bi3+(aq) + MnO 4− (aq)
Oxidation:	4 H 2O(l) + Mn2+(aq) → MnO− 4(aq) + 8 H+(aq) + 5 e−
Reduction:	2 e− + 6 H+ + BiO− 3(s) → Bi3+(aq) + 3 H 2O(l)

The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa):

8 H
2O(l) + 2 Mn2+(aq) → 2 MnO−
4(aq) + 16 H+(aq) + 10 e−

10 e− + 30 H+ + 5 BiO−
3(s) → 5 Bi3+(aq) + 15 H
2O(l)

Adding these two reactions eliminates the electrons terms and yields the balanced reaction:

14 H+(aq) + 2 Mn2+(aq) + 5 NaBiO
3(s) → 7 H
2O(l) + 2 MnO−
4(aq) + 5 Bi3+(aq) + 5 Na+(aq)

Basic media

In basic media, OH⁻ ions and water are added to half reactions to balance the overall reaction.

For example, in the reaction between potassium permanganate and sodium sulfite:

Unbalanced reaction:	KMnO 4 + Na 2SO 3 + H 2O → MnO 2 + Na 2SO 4 + KOH
Reduction:	3 e− + 2 H 2O + MnO 4− → MnO 2 + 4 OH−
Oxidation:	2 OH− + SO 32− → SO 42− + H 2O + 2 e−

Balancing the number of electrons in the two half-cell reactions gives:

6 e− + 4 H
2O + 2 MnO
4⁻ → 2 MnO
2 + 8 OH⁻

6 OH⁻ + 3 SO
3²⁻ → 3 SO
4²⁻ + 3 H
2O + 6 e−

Adding these two half-cell reactions together gives the balanced equation:

2 KMnO
4 + 3 Na
2SO
3 + H
2O → 2 MnO
2 + 3 Na
2SO
4 + 2 KOH

Memory aids

The key terms involved in redox are often confusing to students.^[7][8] For example, an element that is oxidized loses electrons; however, that element is referred to as the reducing agent. Likewise, an element that is reduced gains electrons and is referred to as the oxidizing agent.^[9] Acronyms or mnemonics are commonly used^[10] to help remember what is happening:

• "OIL RIG"—Oxidation Is Loss of electrons, Reduction Is Gain of electrons.^{[7][8][9][10]}

• "LEO the lion says GER" — Loss of Electrons is Oxidation, Gain of Electrons is Reduction.^{[7][8][9][10]}

• "LEORA says GEROA" — Loss of Electrons is Oxidation (Reducing Agent), Gain of Electrons is Reduction (Oxidizing Agent).^[9]

• "RED CAT" and "AN OX", or "AnOx RedCat" ("an ox-red cat"), — Reduction occurs at the Cathode and the Anode is for Oxidation.

• "RED CAT gains what AN OX loses" - Reduction occurs at the Cathode gains (electrons) what Anode Oxidation loses (electrons).

Lewis acids and bases

From Wikipedia, the free encyclopedia

Diagram of Lewis acids and bases

Lewis acid is a chemical species that reacts with a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that donates a pair of electrons to a Lewis acid to form a Lewis adduct. For example, OH⁻ and NH₃ are Lewis bases, because they can donate a lone pair of electrons. In the adduct, the Lewis acid and base share an electron pair furnished by the Lewis base.^[1] Usually the terms Lewis acid and Lewis base are defined within the context of a specific chemical reaction. For example, in the reaction of Me₃B and NH₃ to give Me₃BNH₃, Me₃B acts as a Lewis acid, and NH₃ acts as a Lewis base. Me₃BNH₃ is the Lewis adduct. The terminology refers to the contributions of Gilbert N. Lewis.

Depicting adducts

In many cases, the interaction between the Lewis base and Lewis acid in a complex is indicated by an arrow indicating the Lewis base donating electrons toward the Lewis acid using the notation of a dative bond—for example, Me₃B←NH₃. Some sources indicate the Lewis base with a pair of dots (the explicit electrons being donated), which allows consistent representation of from the base itself to the complex with the acid:

Me₃B +  :NH₃ → Me₃B:NH₃

A center dot may also be used: Me₃B·NH₃, similar to the notation for hydrate coordination in various crystals. In general, however, the donor–acceptor bond is viewed as simply somewhere along a continuum between idealized covalent bonding and ionic bonding.^[2]

Examples

Lewis diagrams showing the formation of the ammonium ion (diagram should show adduct having a positive charge)

Classically, the term "Lewis acid" is restricted to trigonal planar species with an empty p orbital, such as BR₃ where R can be an organic substituent or a halide.^{[citation needed]} For the purposes of discussion, even complex compounds such as Et₃Al₂Cl₃ and AlCl₃ are treated as trigonal planar Lewis acids. Metal ions such as Na⁺, Mg²⁺, and Ce³⁺, which are invariably complexed with additional ligands, are often sources of coordinatively unsaturated derivatives that form Lewis adducts upon reaction with a Lewis base. Other reactions might simply be referred to as "acid-catalyzed" reactions. Some compounds, such as H₂O, are both Lewis acids and Lewis bases, because they can either accept a pair of electrons or donate a pair of electrons, depending upon the reaction.

Lewis acids are diverse. Simplest are those that react directly with the Lewis base. But more common are those that undergo a reaction prior to forming the adduct.

• Examples of Lewis acids based on the general definition of electron pair acceptor include:
  • the proton (H⁺) and acidic compounds onium ions, such as NH₄⁺ and H₃O⁺
  • metal cations, such as Li⁺ and Mg²⁺, often as their aquo or ether complexes,
  • trigonal planar species, such as BF₃ and carbocations H₃C⁺
  • pentahalides of phosphorus, arsenic, and antimony
  • electron poor π-systems, such as enones and tetracyanoethylenes.

Again, the description of a Lewis acid is often used loosely. For example, in solution, bare protons do not exist.

Simple Lewis acids

The most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other compounds exhibit this behavior:

BF₃ + F⁻ → BF₄⁻

In this adduct, all four fluoride centres (or more accurately, ligands) are equivalent.

BF₃ + OMe₂ → BF₃OMe₂

Both BF₄⁻ and BF₃OMe₂ are Lewis base adducts of boron trifluoride.
In many cases, the adducts violate the octet rule, such as the triiodide anion:

I₂ + I⁻ → I₃⁻

The variability of the colors of iodine solutions reflects the variable abilities of the solvent to form adducts with the Lewis acid I₂.

In some cases, the Lewis acids are capable of binding two Lewis bases, a famous example being the formation of hexafluorosilicate:

SiF₄ + 2 F⁻ → SiF₆²⁻

Complex Lewis acids

Most compounds considered to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base. Well known cases are the aluminium trihalides, which are widely viewed as Lewis acids. Aluminium trihalides, unlike the boron trihalides, do not exist in the form AlX₃, but as aggregates and polymers that must be degraded by the Lewis base.^[3] A simpler case is the formation of adducts of borane. Monomeric BH₃ does not exist appreciably, so the adducts of borane are generated by degradation of diborane:

B₂H₆ + 2 H⁻ → 2 BH₄⁻

In this case, an intermediate B₂H₇⁻ can be isolated.

Many metal complexes serve as Lewis acids, but usually only after dissociating a more weakly bound Lewis base, often water.

[Mg(H₂O)₆]²⁺ + 6 NH₃ → [Mg(NH₃)₆]²⁺ + 6 H₂O

H⁺ as Lewis acid

The proton (H⁺) ^[4] is one of the strongest but is also one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is heavily solvated (bound to solvent). With this simplification in mind, acid-base reactions can be viewed as the formation of adducts:

• H⁺ + NH₃ → NH₄⁺
• H⁺ + OH⁻ → H₂O

Applications of Lewis acids

Typical example of a Lewis acid in action is in the Friedel–Crafts alkylation reaction.^[2] The key step is the acceptance by AlCl₃ of a chloride ion lone-pair, forming AlCl₄⁻ and creating the strongly acidic, that is, electrophilic, carbonium ion.

RCl +AlCl₃ → R⁺ + AlCl₄⁻

Lewis Bases

A Lewis base is an atomic or molecular species where the HOMO is highly localized. Typical Lewis bases are conventional amines such as ammonia and alkyl amines. Other common Lewis bases include pyridine and its derivatives. Some of the main classes of Lewis bases are

• amines of the formula NH_3−xR_x where R = alkyl or aryl. Related to these are pyridine and its derivatives.
• phosphines of the formula PR_3−xA_x, where R = alkyl, A = aryl.
• compounds of O, S, Se and Te in oxidation state 2, including water, ethers, ketones

The most common Lewis bases are anions. The strength of Lewis basicity correlates with the pK_a of the parent acid: acids with high pK_a's give good Lewis bases. As usual, a weaker acid has a stronger conjugate base.

• Examples of Lewis bases based on the general definition of electron pair donor include:
  • simple anions, such as H⁻ and F⁻.
  • other lone-pair-containing species, such as H₂O, NH₃, HO⁻, and CH₃⁻
  • complex anions, such as sulfate
  • electron rich π-system Lewis bases, such as ethyne, ethene, and benzene

The strength of Lewis bases have been evaluated for various Lewis acids, such as I₂, SbCl₅, and BF₃.^[5]

Heats of binding of various bases to BF3
Lewis base	donor atom	Enthalpy of Complexation (kJ/mol)
Et3N	N	135
quinuclidine	N	150
pyridine	N	128
Acetonitrile	N	60
Et2O	O	78.8
THF	O	90.4
acetone	O	76.0
EtOAc	O	75.5
DMA	O	112
DMSO	O	105
Tetrahydrothiophene	S	51.6
PMe3	P	97.3

Applications of Lewis bases

Nearly all electron pair donors that form compounds by binding transition elements can be viewed as a collections of the Lewis bases – or ligands. Thus a large application of Lewis bases is to modify the activity and selectivity of metal catalysts. Chiral Lewis bases thus confer chirality on a catalyst, enabling asymmetric catalysis, which is useful for the production of pharmaceuticals.
Many Lewis bases are "multidentate," that is they can form several bonds to the Lewis acid. These multidentate Lewis acids are called chelating agents.

Hard and soft classification

Lewis acids and bases are commonly classified according to their hardness or softness. In this context hard implies small and nonpolarizable and soft indicates larger atoms that are more polarizable.

• typical hard acids: H⁺, alkali/alkaline earth metal cations, boranes, Zn²⁺
• typical soft acids: Ag⁺, Mo(0), Ni(0), Pt²⁺
• typical hard bases: ammonia and amines, water, carboxylates, fluoride and chloride
• typical soft bases: organophosphines, thioethers, carbon monoxide, iodide

For example, an amine will displace phosphine from the adduct with the acid BF₃. In the same way, bases could be classified. For example, bases donating a lone pair from an oxygen atom are harder than bases donating through a nitrogen atom. Although the classification was never quantified it proved to be very useful in predicting the strength of adduct formation, using the key concepts

• hard acid — hard base interactions are stronger than hard acid — soft base or soft acid — hard base interactions.
• soft acid — soft base interactions are stronger than soft acid — hard base or hard acid — soft base interactions.

Later investigation of the thermodynamics of the interaction suggested that hard—hard interactions are enthalpy favored, whereas soft—soft are entropy favored.

History

MO diagram depicting the formation of a dative covalent bond between two atoms.

The concept originated with Gilbert N. Lewis who studied chemical bonding. In 1923, Lewis wrote An acid substance is one which can employ an electron lone pair from another molecule in completing the stable group of one of its own atoms.^[6][7] The Brønsted–Lowry acid–base theory was published in the same year. The two theories are distinct but complementary. A Lewis base is also a Brønsted–Lowry base, but a Lewis acid doesn't need to be a Brønsted–Lowry acid. The classification into hard and soft acids and bases (HSAB theory) followed in 1963. The strength of Lewis acid-base interactions, as measured by the standard enthalpy of formation of an adduct can be predicted by the Drago–Wayland two-parameter equation.

Reformulation of Lewis Theory

Lewis had suggested in 1916 that two atoms are held together in a chemical bond by sharing a pair of electrons. When each atom contributed one electron to the bond it was called a covalent bond. When both electrons come from one of the atoms it was called a dative covalent bond or coordinate bond. The distinction is not very clear-cut. For example, in the formation of an ammonium ion from ammonia and hydrogen the ammonia molecule donates a pair of electrons to the proton;^[4] the identity of the electrons is lost in the ammonium ion that is formed. Nevertheless, Lewis suggested that an electron-pair donor be classified as a base and an electron-pair acceptor be classified as acid.
A more modern definition of a Lewis acid is an atomic or molecular species with a localized empty atomic or molecular orbital of low energy. This lowest energy molecular orbital (LUMO) can accommodate a pair of electrons.

Comparison with Brønsted–Lowry Theory

A Lewis base is often a Brønsted–Lowry base as it can donate a pair of electrons to H⁺;^[4] the proton is a Lewis acid as it can accept a pair of electrons. The conjugate base of a Brønsted–Lowry acid is also a Lewis base as loss of H⁺ from the acid leaves those electrons which were used for the A—H bond as a lone pair on the conjugate base. However, a Lewis base can be very difficult to protonate, yet still react with a Lewis acid. For example, carbon monoxide is a very weak Brønsted–Lowry base but it forms a strong adduct with BF₃.

In another comparison of Lewis and Brønsted–Lowry acidity by Brown and Kanner,^[8] 2,6-di-t-butylpyridine reacts to form the hydrochloride salt with HCl but does not react with BF₃. This example demonstrates that steric factors, in addition to electron configuration factors, play a role in determining the strength of the interaction between the bulky di-t-butylpyridine and tiny proton.
A Brønsted–Lowry acid is a proton donor, not an electron-pair acceptor.