Mix bromomethane, CH₃Br, with hydroxide ion and you get methanol, CH₃OH. Hydroxide uses one of its lone pairs to make a new bond to the carbon, and bromide ion leaves. Hydroxide brought the electron pair, so it is the nucleophile (Nucleophiles). The carbon received the pair, so it is the electrophile.
Every reaction step in this course that moves an electron pair brings together a nucleophile and an electrophile. The curved arrow starts at the nucleophile's electrons and points at the electrophile's electron-poor atom. Once you have found both, you know where the new bond will form. Almost every mechanism in the course repeats that one move with different molecules, which is why reading a molecule for its electron density is worth more than memorizing reactions.
What it is
An electrophile is a species that accepts an electron pair to form a new bond. The name means "electron-loving," but electron-poor is the more useful description. An atom can be electron-poor in any of three ways:
- a partial positive charge, δ+, left behind when a more electronegative neighbor pulls a shared pair toward itself (Bond polarity);
- a full positive charge, a formal charge of +1;
- an empty orbital, with room for an incoming pair.
You have met this definition before. In Acids & Bases, an electron-pair acceptor was called a Lewis acid. The two words describe the same behavior: "Lewis acid" is the acid–base name for it, and "electrophile" is the name used when drawing mechanisms.
A full positive charge is not required, and most of the time it is absent. The great majority of electrophilic carbons you will meet are neutral atoms carrying a δ+ from a neighbor.
Acetone is (CH₃)₂C=O. Oxygen is much more electronegative than carbon (3.44 against 2.55), so it pulls the C=O bonding electrons toward itself. The carbonyl carbon is left δ+ and the oxygen δ−.
So the carbon is the electrophilic atom, the site a nucleophile attacks. The oxygen, with its two lone pairs, looks like the more interesting atom, but it is electron-rich. Acids protonate it; nucleophiles do not attack it.
Both atoms have a formal charge of zero. The δ+ comes from electronegativity alone, which is why formal charges by themselves will not find an electrophile for you.
The main families
A δ+ carbon next to a leaving group. Go back to bromomethane. Bromine (2.96) is more electronegative than carbon (2.55), so the carbon of the C–Br bond is δ+. In CH₃–Br it is the carbon that is electrophilic, not the bromine, even though bromine is the atom that eventually leaves.
A leaving group is the atom or group that departs as the nucleophile bonds, taking the old bond's electron pair with it. In bromomethane it is bromine, which leaves as bromide ion, Br⁻. Chlorine and iodine on carbon behave the same way, and so does the tosylate group, OTs, described under "Making an electrophile stronger" below. What makes a group leave easily is the subject of Leaving groups, later in this chapter.
A carbonyl carbon. Acetone, above, is the example of this family. The carbonyl carbon of every aldehyde, ketone, ester, amide and acid chloride is δ+ in the same way, and Carbonyl Chemistry is built on it.
A carbocation. A carbocation is a carbon with only three bonds, no lone pair and a positive charge. The formal-charge count gives 4 − 0 − 3 = +1. One forms when a leaving group departs from a carbon, taking the whole bonding pair with it, before any nucleophile has arrived. The carbon is sp² and flat, and its empty p orbital stands at right angles to that plane, one lobe above it and one below.
A full positive charge together with an empty orbital makes carbocations some of the strongest electrophiles in the course. They appear as short-lived intermediates in the SN1 and E1 reactions of Substitution & Elimination and in many additions to alkenes. Carbocations, later in this chapter, covers their shape and what makes one more stable than another.
The proton, H⁺. A proton is the simplest electrophile there is. It has no electrons at all, only an empty 1s orbital, and it takes the nearest available lone pair or π bond. That is the whole of the Brønsted acid–base chemistry in Acids & Bases, and it is the first step of a great many mechanisms. In solution a bare proton never floats free. An acid such as H₃O⁺ or H–Br carries it and hands it over.
A polarized reagent. Some reagents have an electrophilic atom bonded to a more electronegative one. In H–Br, bromine pulls the shared pair toward itself and leaves the hydrogen δ+. A nucleophile attacks the H end, and bromide leaves with the pair. This is how H–Br adds to an alkene: the alkene's π electrons take the H (Alkenes & Alkynes).
Lewis acids. In BF₃, boron has only six electrons around it (three bonding pairs) and an empty p orbital. AlCl₃ is built the same way. Metal salts such as FeBr₃ and ZnCl₂ are Lewis acids too. They accept a pair by taking on one more group, not by filling an empty p orbital; the Lewis acids section of Acids & Bases shows how.
Induced-dipole electrophiles. Even a nonpolar molecule can become electrophilic when a nucleophile comes close. Br₂ is two identical atoms, so its bond has no δ+ end. As Br₂ approaches an alkene end-on, the alkene's π electrons repel the Br–Br pair toward the far bromine. The near bromine is left with a temporary δ+, and that is the atom the alkene attacks (Alkenes & Alkynes).
Making an electrophile stronger
Take acetone again. Its carbonyl carbon carries only a partial positive charge. Add a strong acid, and one of the oxygen's lone pairs bonds to a proton. The product is a cation with two resonance structures. In one, the + sits on oxygen and the C=O stays intact. In the other, the C=O π pair has moved onto oxygen, and the + sits on carbon. The actual ion is a blend of the two, so its carbon carries a sizeable share of a full positive charge. That makes it far more electrophilic than the carbon of neutral acetone.
This is protonation, the most common way to strengthen an electrophile, and it is why so many carbonyl reactions are run with an acid catalyst.
Coordination to a Lewis acid does the same job without a proton. The oxygen shares one of its lone pairs with a Lewis acid such as BF₃ and becomes positive, so it pulls harder on the carbon. This route works on molecules that a strong acid would destroy.
Converting to a better leaving group is the third route. Ethanol's C–O carbon is δ+, yet nothing happens there, because hydroxide, HO⁻, does not leave. Now swap the H of the OH for a tosyl group. Ts is short for tosyl, p-toluenesulfonyl, CH₃C₆H₄SO₂–: a sulfur carrying two S=O oxygens and a benzene ring, with a methyl group on the ring carbon opposite the sulfur. The product, CH₃CH₂–OTs, is ethyl tosylate, and its leaving group, OTs, is called tosylate.
The carbon's δ+ barely changes. What changes is the group on it: tosylate leaves easily, so the carbon now reacts with nucleophiles that ignore ethanol. Leaving groups explains why.
Ranking electrophiles
Exam questions rarely ask whether something is an electrophile. They ask which of two is the better one. Two rules cover almost every comparison at this stage, and both rest on induction and resonance, the two effects used to rank acids in Acids & Bases. Induction is a pull on electrons through the σ bonds; resonance moves electrons through a π system and is drawn with curved arrows.
Compare acetaldehyde, CH₃–CHO, with trifluoroacetaldehyde, CF₃–CHO. The three fluorines pull electron density out of the carbonyl carbon by induction, which deepens its δ+. CF₃–CHO is a far stronger electrophile than CH₃–CHO. So a neighbor that pulls density away makes a carbon more electrophilic. A halogen, a second carbonyl group or a positively charged atom next door all work this way.
Now compare acetaldehyde with acetone. Acetaldehyde's carbonyl carbon carries one CH₃ and an H; acetone's carries two CH₃ groups. An alkyl group such as CH₃ donates a little electron density, which partly cancels the carbon's δ+. It also takes up more room than a hydrogen and gets in the nucleophile's way. So a ketone is less electrophilic, and harder to reach, than an aldehyde. Aldehydes & ketones puts numbers on the difference.
A lone pair next door donates far more. In acetamide, CH₃CONH₂, the nitrogen's lone pair is shared with the carbonyl carbon by resonance. One resonance structure has a C=N double bond in place of the C–N single bond, with a + on nitrogen and a − on oxygen. The carbon is left much less δ+, which is why an amide's carbonyl carbon is so unreactive.
A full negative charge next door donates hardest of all. In a carboxylate ion, RCOO⁻, the carbonyl carbon sits beside a negatively charged oxygen, and it is a very poor electrophile. So a neighbor that gives density back makes a carbon less electrophilic, and a lone pair or a negative charge gives far more than an alkyl group does.
Chlorine both pulls and gives. It pulls by induction, and its lone pairs give back by resonance much less than an ester oxygen's do. That is why an acid chloride, RCOCl, is a much better electrophile than an ester, RCOOR′.
Put the two rules together and any set of carbonyl compounds sorts itself: the more a neighbor takes, the better the electrophile; the more it gives, the worse. The full ordering of the acid derivatives belongs to Carboxylic Acids & Derivatives, but the reasons are the ones on this page.
Finding every site in one molecule
Ranking compares two molecules. A different question asks where, inside one molecule, a nucleophile will attack. Answer it by walking the molecule carbon by carbon and asking two things at each one. Is it δ+? And does it have a way to accept a new pair: a group that can leave, or a π bond that can break?
Take 4-chlorobutan-2-one, ClCH₂–CH₂–CO–CH₃, and ask which carbon a nucleophile attacks. Number the carbons as the name does: C1 is the CH₃ and C4 carries the chlorine. Each of the four needs an argument, not a glance.
Carbon 1, the CH₃. Apart from the carbonyl carbon, it is bonded only to hydrogens. Nothing pulls density from it, and nothing on it can leave. It is not a target.
Carbon 3, the middle CH₂. It sits between the CH₂Cl and the carbonyl group, so both pull on it by induction. Induction fades quickly with distance, though, and the chlorine and the oxygen are each two bonds away. Nothing on this carbon can leave either. It is not a target.
Carbon 4, the CH₂Cl. The chlorine makes it clearly δ+, and chloride can leave. This is a real electrophilic site. A nucleophile that prefers a saturated carbon (one with four single bonds), such as iodide, replaces the chloride here, in an SN2 reaction (Substitution & Elimination).
Carbon 2, the carbonyl carbon. Its double bond to oxygen makes it the most strongly δ+ carbon in the molecule. When a nucleophile bonds to it, the C=O π bond breaks and sends its pair onto oxygen, which, being electronegative, holds the extra pair well. The carbon still ends up with four bonds.
The answer: there are two sites. The carbonyl carbon is the more electrophilic, and a hydride donor (a reagent that delivers H⁻) adds there. The CH₂Cl carbon is the second site, and a nucleophile such as iodide goes there. A molecule with two electrophilic sites is normal; the question is which one a given nucleophile picks, not whether a site exists.
Keep the difference between the two sites in mind. Carbon 4 reacts only if chloride leaves, so how good a leaving group chloride is sets its rate. Carbon 2 needs nothing to leave at all. That is why substitution chemistry and carbonyl chemistry look so different, even though both begin with a nucleophile attacking a δ+ carbon.
Electrophilicity is relative
Nothing is an electrophile or a nucleophile in the abstract. Both words name a role in a particular reaction, and the same molecule can play either role.
Water is the simplest case. Its oxygen has lone pairs, and toward a carbocation water acts as a nucleophile. Its hydrogens are δ+, and a strong base such as the amide ion, NH₂⁻, attacks one of them. In that reaction water is the electrophile.
An alkene is a nucleophile toward Br₂. Yet an alkene becomes electrophilic once it is conjugated to a carbonyl group, meaning its C=C is joined directly to the C=O carbon. In but-3-en-2-one, CH₂=CH–CO–CH₃, the carbon bonded to the carbonyl carbon is the alpha (α) carbon. The next one along, at the far end of the double bond, is the beta (β) carbon. One resonance structure puts a + on the β carbon, so that carbon is electron-poor. The next section, Electron-rich vs. electron-poor atoms, looks more closely at resonance that reaches along a chain like this.
So the question is never "is this molecule an electrophile?" It is "in this pairing, which side is electron-poor?"
What carries forward
The δ+ carbon that carries a leaving group drives all of Substitution & Elimination. The electrophilic carbonyl carbon drives Carbonyl Chemistry, Carboxylic Acids & Derivatives and Enolate Chemistry. Electrophilic aromatic substitution in Aromatic Chemistry is named for this idea outright. Making an unusually strong electrophile comes up every time a reaction needs a push. Protonated carbonyls are one example from this page. Others come later: acylium ions, R–C≡O⁺, and bromonium ions, in which one positive bromine bridges both carbons of a former alkene.