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The Study of Organic Chemistry

Organic chemistry is primarily devoted to the unique properties of the carbon atom and its compounds. These compounds play a critical role in biology and ecology, Earth sciences and geology, physics, industry, medicine and — of course — chemistry. At first glance, the new material that organic chemistry brings to the table may seem complicated and daunting, but all it takes is concentration and perseverance. Millions of students before you have successfully passed this course and you can too!

This field of chemistry is based less on formulas and more on reactions between various molecules under different conditions. Whereas a typical general chemistry question may ask a student to compute an answer with an equation from the chapter that they memorized, a more typical organic chemistry question is along the lines of "what product will form when substance X is treated with solution Y and bombarded by light". The key to learning organic chemistry is to understand it rather than cram it in the night before a test. It is all well and good to memorize the mechanism of Michael addition, but a superior accomplishment would be the ability to explain why such a reaction would take place.

As in all things, it is easier to build up a body of new knowledge on a foundation of solid prior knowledge. Students will be well served by much of the knowledge brought to this subject from the subject of General chemistry. Concepts with particular importance to organic chemists are covalent bonding, Molecular Orbit theory, VSEPR Modeling, understanding acid/base chemistry vis-a-vis pKa values, and even trends of the periodic table. This is by no means a comprehensive list of the knowledge you should have gained already in order to fully understand the subject of organic chemistry, but it should give you some idea of the things you need to know to succeed in an organic chemistry test or course.

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Copyright

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Authors

The authors of this book are:

1. Karl Wick Citizen of the United States of America, and living there

   The initiator of this project is Karl Wick, who is finishing up his premed science courses at the Cleveland State University in Cleveland, Ohio. At the time of this writing (7/15/03) I have been the sole contributor but as time goes on it will become a group project "of the people" as many contribute and improve it by bits and pieces.

2. Justin Johnson Citizen of the United States of America, residing therein, born 1975

   Justin Johnson(JSJohnson) is a pre-medical undergraduate student at Indiana University Purdue University Indianapolis. He first read this book in the summer of 2005, and began contributing to it in the spring of 2006.

3. David Rose

   David Rose (Ghostal) is a chemical engineering undergraduate at Michigan Technological University. He began adding to this book in October, 2004.

4. #Patrick Holder

   Patrick Holder (Lineweaver) is a PhD Graduate Student working for Matt Francis at UC-Berkeley in the Department of Chemistry. He began adding to this book in December, 2005.

5. Zachary T. Tackett, Zach is a student at Marshall University majoring in bio-organic chemistry.
6. Spammy
7. Igoroisha
8. Jack H
9. Avijay86
10. xcentaur
11. Goh Liang Song Citizen of the Republic of Singapore

    Goh Liang Song (User:Gohliangsong) Citizen of the Republic of Singapore. He graduated from the National University of Singapore, majoring in Chemistry. He began adding to this book in August, 2006.

12. Pete Davis Citizen of the United States of America. I joined this project in November 2006. There are few of us working on it at the moment. Feel free to join in.
13. Ewen McLaughlin UK Citizen, living in Wales. I started adding bits here in February 2006. I'm a teacher, so I suppose I might be more help with how to present information than with the information itself.
14. Shalom

And many anonymous Wikibook contributors.

Many thanks to Jimbo Wales for paying for the bandwidth and for the many other ways he has been a great support to this project at every step.

Foreword

Purpose and mission

This book should become the gold standard of organic chemistry texts in the areas of accuracy, usability, flexibility, and connection with its audience. As this text is developed it will always be available online, be printable, and freely distributable. This text should eliminate all or much of the cost for owning an up-to-the-minute, top-quality college-level organic chemistry text, as it and all its derivative works will remain free: free as in speech as well as free as in beer. Although you could pay for a printed version if you wanted to.

Content and Contributions

This is, to the best of our knowlege, the world's first and only open content organic chemistry textbook. Its users will tweak and refine this book until there is no better book. We are confident that this will happen because the process has already been seen to work many times on the Wikipedia site.

All content contained herein is available under licences that allow free distribution. You can copy it, print it, sell it, and create derivative works from it.

Our restriction: if you create derivative works, make them available to others in a way that they can easily copy and distribute them, as we have done for you.

We link to some pages outside our server. Any of this content not found under the Wikipedia site and subsites is not ensured to be under the same license; it in fact is most likely not.

Navigation

I like the navigation that I have worked out in the Foundational concepts of organic chemistry page and immediate subpages with links at the top of the page to the immediate next and previous pages and the various levels, and links at the bottom to the next and previous chapters, etc. If you want to help out please help me get all of the pages in the same format.

Licensing

All work in this book is released at the moment only under the GNU FDL license. However this is only one of many similar open content licenses, and may not be the license of choice for everyone. To take content written by Karl Wick from this book for release under other licenses please contact the author through this page's associated talk page.

How to study organic chemistry

One of the main difficulties students have with organic chemistry is organizing the information in their minds. By the second semester of organic chemistry, students will learn over 100 chemical reactions. Consequently, it is vital that students take time to not only organize the information, but also to understand it. Indeed, excellent organic professors will tell you, contrary to popular belief, that you do not really need to memorize anything for organic chemistry, instead you simply need to understand it. By truly learning something, rather than memorizing it, you will be able to apply concepts beyond what you are memorizing.

When you see something in the textbook, always ask why something is the case. Do research, try to find out the answer. By taking this approach you will enrich your learning experience, and the information will be "locked" in your mind.

Each person may have a slightly different method that helps him or her learn organic chemistry the quickest and with least pain. The basic rule of thumb is to use a method that you find most helpful and stick with it. Various study methods include flash cards, molecular model kits, group study, writing chemical reactions on blackboards, others just take the class over and over until they "get it".

The writers would recommend to buy a molecular model kit so you can hold in your hand and visualize in your mind how the molecules look in three-dimensional space. If you can't get access to models or can't afford them, look online for sites that use the Jmol application or other rendering software that allow you to virtually rotate molecules.

It cannot be stressed enough that you must be able to visualize molecules in organic chemistry. The 3 dimensional structure of molecules often plays a crucial part in reactions. It can be the deciding factor in whether a reaction even happens, it can decide how fast it happens, and it can decide what the product(s) of the reaction is going to be. If you can't visualize the 3D structure, you won't be able to understand what's happening.

Sports analogy

You can think of the different elements and functional groups as players in a game and the organic reactions as the plays. Just as each player or team has different strengths or characteristics and uses strategies to achieve what they want, organic chemists use the properties of each chemical to play off the others in order to achieve a desired end result.

Language analogy

You can also think of organic chemistry like learning a foreign language. The atoms, for example, carbon and hydrogen and oxygen and nitrogen, are the letters of the alphabet. The structural theory of organic chemistry, namely, the tetravalencey of carbon, may be considered the essential underlying grammatical rule. All organic compounds are assembled under these grammatical rules, and may be considered words. The reactions of organic compounds may be perceived as the assembly of these words into sentences. A language analogy is also useful at this point, because the grammatical rules that control the assembly of sentences (formation of the products of organic reactions!) may be found in the study of organic reaction mechanisms.

Therefore, it is not necessary to memorize individual reactions. Overall patterns of reactivity become obvious when the mechanism of the reaction is investigated. Moreover, like any language, you have to practice it constantly. The more you "read" and "speak" chemical reactions and understand the mechanisms by which they proceed, the more fluent you will become. When you finish organic chemistry, you will literally be able to read, write, and speak in a foreign language. However, it is important to note that the language of organic chemistry is far simpler than any language people use for general communication! The words mean exactly what they mean, and the basic rules almost never change. But organic chemistry is far from a dead science. In fact, it is one of the most active and rapidly advancing areas in modern science today.

Research produces new knowledge, and the potential to formulate new rules. Perhaps you will make some of these discoveries, and future students will refer to your rules.

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Unit 1: Foundational concepts of organic chemistry

History of organic chemistry

Brief History

Jöns Jacob Berzelius, a physician by trade, first coined the term "organic chemistry" in 1807 for the study of compounds derived from biological sources. Up through the early 19th century, naturalists and scientists observed critical differences between compounds that were derived from living things and those that were not.

Chemists of the period noted that there seemed to be an essential yet inexplicable difference between the properties of the two different types of compounds. The vital force theory (sometimes called "vitalism") was therefore proposed (and widely accepted) as a way to explain these differences. Vitalism proposed that there was a something called a "vital force" which existed within organic material but did not exist in any inorganic materials.

Synthesis of Urea

Friedrich Wöhler is widely regarded as a pioneer in organic chemistry as a result of his synthesizing of the biological compound urea (a component of urine in many animals) utilizing what is now called "the Wöhler synthesis."

Wöhler mixed silver or lead cyanate with ammonium nitrate; this was supposed to yield ammonium cyanate as a result of an exchange reaction, according to Berzelius's dualism theory. Wöhler, however, discovered that the end product of this reaction is not ammonium cyanate (NH₄OCN), an inorganic salt, but urea ((NH₂)₂CO), a biological compound. (Furthermore, heating ammonium cyanate turns it into urea.) Faced with this result, Berzelius had to concede that (NH₂)₂CO and NH₄OCN were isomers. Until this discovery in the year 1828, it was widely believed by chemists that organic substances could only be formed under the influence of the "vital force" in the bodies of animals and plants. Wöhler's synthesis dramatically proved that view to be false.

Urea synthesis was a critical discovery for biochemists because it showed that a compound known to be produced in nature only by biological organisms could be produced in a laboratory under controlled conditions from inanimate matter. This "in vitro" synthesis of organic matter disproved the common theory (vitalism) about the vis vitalis, a transcendent "life force" needed for producing organic compounds.

Organic vs Inorganic Chemistry

Although originally defined as the chemistry of biological molecules, organic chemistry has since been redefined to refer specifically to carbon compounds — even those with non-biological origin. Some carbon molecules are not considered organic, with carbon dioxide being the most well known and most common inorganic carbon compound, but such molecules are the exception and not the rule.

Organic chemistry focuses on carbon and following movement of the electrons in carbon chains and rings, and also how electrons are shared with other carbon atoms and heteroatoms. Organic chemistry is primarily concerned with the properties of covalent bonds and non-metallic elements, though ions and metals do play critical roles in some reactions.

The applications of organic chemistry are myriad, and include all sorts of plastics, dyes, flavorings, scents, detergents, explosives, fuels and many, many other products. Read the ingredient list for almost any kind of food that you eat — or even your shampoo bottle — and you will see the handiwork of organic chemists listed there.

Major Advances in the Field of Organic Chemistry

Of course no description of a text should be without at least a mention of Antoine Laurent Lavoisier. The French chemist is often called the "Father of Modern Chemistry" and his place is first in any pantheon of great chemistry figures. Your general chemistry textbook should contain information on the specific work and discoveries of Lavoisier — they will not be repeated here because his discoveries did not relate directly to organic chemistry in particular.

Berzelius and Wöhler are discussed above, and their work was foundational to the specific field of organic chemistry. After those two, three more scientists are famed for independently proposing the elements of structural theory. Those chemists were August Kekulé, Archibald Couper and Alexander Butlerov.

Kekulé was a German, an architect by training, and he was perhaps the first to propose that the concept of isomerism was due to carbon's proclivity towards forming four bonds. Its ability to bond with up to four other atoms made it ideal for forming long chains of atoms in a single molecule, and also made it possible for the same number of atoms to be connected in an enormous variety of ways. Couper, a Scot, and Butlerov, a Russian, came to many of the same conclusions at the same time or just a short time after.

Through the nineteenth century and into the twentieth, experimental results brought to light much new knowledge about atoms, molecules and molecular bonding. In 1916 it was Gilbert Lewis of U.C. Berkeley who described covalent bonding largely as we know it today (electron-sharing). Nobel laureate Linus Pauling further developed Lewis' concepts by proposing resonance while he was at the California Institute of Technology. At about the same time, Sir Robert Robinson of Oxford University focused primarily on the electrons of atoms as the engines of molecular change. Sir Christopher Ingold of University College, London, organized what was known of organic chemical reactions by arranging them in schemes we now know of as mechanisms, in order to better understand the sequence of changes in a synthesis or reaction.

The field of organic chemistry is probably the most active and important field of chemistry at the moment, due to its extreme applicability to both biochemistry (especially in the pharmaceutical industry) and petrochemistry (especially in the energy industry). Organic chemistry has a relatively recent history, but it will have an enormously important future, affecting the lives of everyone around the world for many, many years to come.

Atomic structure

A simple model of a lithium atom. Not to scale!

Atomic Structure

Atoms are made up of a nucleus and electrons that orbit the nucleus. The nucleus consists of protons and neutrons. An atom in its natural, uncharged state has the same number of electrons as protons.

The nucleus

The nucleus is made up of protons, which are positively charged, and neutrons, which have no charge. Neutrons and protons have about the same mass, and together account for most of the mass of the atom. Each of these particles is made up of even smaller particles, though the existence of these particles does not come into play at the energies and time spans in which most chemical reactions occur. The ratio of protons to neutrons is fairly critical, and any departure from the optimum range will lead to nuclear instability and thus radioactivity.

Electrons

The electrons are negatively charged particles. The mass of an electron is about 2000 times smaller than that of a proton or neutron at 0.00055 amu. Electrons circle so fast that it cannot be determined where electrons are at any point in time, rather, we talk about the probability of finding an electron at a point in space relative to a nucleus at any point in time. The image depicts the old Bohr model of the atom, in which the electrons inhabit discrete "orbitals" around the nucleus much like planets orbit the sun. This model is outdated. Current models of the atomic structure hold that electrons occupy fuzzy clouds around the nucleus of specific shapes, some spherical, some dumbbell shaped, some with even more complex shapes. Even though the simpler Bohr model of atomic structure has been superseded, we still refer to these electron clouds as "orbitals". The number of electrons and the nature of the orbitals they occupy basically determines the chemical properties and reactivity of all atoms and molecules.

Shells and Orbitals

Electron orbitals

Electrons orbit atoms in clouds of distinct shapes and sizes. The electron clouds are layered one inside the other into units called shells (think nested Russian dolls), with the electrons occupying the simplest orbitals in the innermost shell having the lowest energy state and the electrons in the most complex orbitals in the outermost shell having the highest energy state. The higher the energy state, the more energy the electron has, just like a rock at the top of a hill has more potential energy than a rock at the bottom of a valley. The main reason why electrons exist in higher energy orbitals is because only two electrons can exist in any orbital. So electrons fill up orbitals, always taking the lowest energy orbitals available. An electron can also be pushed to a higher energy orbital, for example by a photon. Typically this is not a stable state and after a while the electron descends to lower energy states by emitting a photon spontaneously. These concepts will be important in understanding later concepts like optical activity of chiral compounds as well as many interesting phenomena outside the realm of organic chemistry (for example, how lasers work).

Wave nature of electrons

The result of this observation is that electrons are not just in simple orbit around the nucleus as we imagine the moon to circle the earth, but instead occupy space as if they were a wave on the surface of a sphere.

If you jump a jumprope you could imagine that the wave in the rope is in its fundamental frequency. The high and low points fall right in the middle, and the places where the rope doesn't move much (the nodes) occur only at the two ends. If you shake the rope fast enough in a rythmic way, using more energy than you would just jumping rope, you might be able to make the rope vibrate with a wavelength shorter than the fundamental. You then might see that the rope has more than one place along its length where it vibrates from its highest spot to its lowest spot. Furthermore, you'll see that there are one or more places (or nodes) along its length where the rope seems to move very little, if at all.

Or consider stringed musical instruments. The sound made by these instruments comes from the different ways, or modes the strings can vibrate. We can refer to these different patterns or modes of vibrations as linear harmonics. Going from there, we can recognize that a drum makes sound by vibrations that occur across the 2-dimensional surface of the drumhead. Extending this now into three dimensions, we think of the electron as vibrating across a 3-dimensional sphere, and the patterns or modes of vibration are referred to as spherical harmonics. The mathematical analysis of spherical harmonics were worked out by the French mathematician Legendre long before anyone started to think about the shapes of electron orbitals. The algebraic expressions he developed, known as Legendre polynomials, describe the three dimension shapes of electron orbitals in much the same way that the expression x²+y² = z² describes a circle (or, for that matter, a drumhead). Many organic chemists need never actually work with these equations, but it helps to understand where the pictures we use to think about the shapes of these orbitals come from.

Electron shells

Each different shell is subdivided into one or more orbitals, which also have different energy levels, although the energy difference between orbitals is less than the energy difference between shells.

Longer wavelengths have less energy; the s orbital has the longest wavelength allowed for an electron orbiting a nucleus and this orbital is observed to have the lowest energy.

Each orbital has a characteristic shape which shows where electrons most often exist. The orbitals are named using letters of the alphabet. In order of increasing energy the orbitals are: s, p, d, and f orbitals.

As one progresses up through the shells (represented by the principal quantum number n) more types of orbitals become possible. The shells are designated by numbers. So the 2s orbital refers to the s orbital in the second shell.

S orbital

The s orbital is the orbital lowest in energy and is spherical in shape. Electrons in this orbital are in their fundamental frequency. This orbital can hold a maximum of two electrons.

P orbital

The next lowest-energy orbital is the p orbital. Its shape is often described as like that of a dumbbell. There are three p-orbitals each oriented along one of the 3-dimensional coordinates x, y or z. Each of these three "p" orbitals can hold a maximum of two electrons.

These three different p orbitals can be referred to as the p_x, p_y, and p_z.

The s and p orbitals are important for understanding most of organic chemistry as these are the orbitals that are occupied by the type of atoms that are most common in organic compounds.

D and F orbitals

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There are also D and F orbitals. D orbitals are present in transition metals. Sulfur and phosphorus have empty D orbitals. Compounds involving atoms with D orbitals do come into play, but are rarely part of an organic molecule. F are present in the elements of the lanthanide and actinide series. Lanthanides and actinides are mostly irrelevant to organic chemistry.

Filling electron shells

When an atom or ion receives electrons into its orbitals, the orbitals and shells fill up in a particular manner.

There are three principles that govern this process:

1. the Pauli exclusion principle,
2. the Aufbau (build-up) principle, and
3. Hund's rule.

Pauli exclusion principle

No more than one electron can have all four quantum numbers the same. What this translates to in terms of our pictures of orbitals is that each orbital can only hold two electrons, one "spin up" and one "spin down".

The Pauli exclusion principle is a quantum mechanical principle formulated by Wolfgang Pauli in 1925, which states that no two identical fermions may occupy the same quantum state simultaneously. It is one of the most important principles in physics, primarily because the three types of particles from which ordinary matter is made—electrons, protons, and neutrons—are all subject to it. The Pauli exclusion principle underlies many of the characteristic properties of matter, from the large-scale stability of matter to the existence of the periodic table of the elements.

Pauli exclusion principle follows mathematically from definition of wave function for a system of identical particles - it can be either symmetric or antisymmetric (depending on particles' spin).

Particles with antisymmetric wave function are called fermions - they have to obey the Pauli exclusion principle. Apart from the familiar electron, proton and neutron, these include the neutrinos, the quarks (from which protons and neutrons are made), as well as some atoms like helium-3. All fermions possess "half-integer spin", meaning that they possess an intrinsic angular momentum whose value is given by Dirac's constant (Planck's constant divided by 2π) times a half-integer (1/2, 3/2, 5/2, etc.). In the theory of quantum mechanics, fermions are described by "antisymmetric states", which are explained in greater detail in the article on identical particles.

Particles with integer spin have symmetric wave function and are called bosons, in contrast to fermions they share same quantum states. Examples of bosons include the photon and the W and Z bosons.

Build-up principle

According to the principle, electrons fill orbitals starting at the lowest available energy states before filling higher states (e.g. 1s before 2s).

You may consider an atom as being "built up" from a naked nucleus by gradually adding to it one electron after another, until all the electrons it will hold have been added. Much as one fills up a container with liquid from the bottom up, so also are the orbitals of an atom filled from the lowest energy orbitals to the highest energy orbitals.

However, the three p orbitals of a given shell all occur at the same energy level. So, how are they filled up? Is one of them filled full with the two electrons it can hold first, or do each of the three orbitals receive one electron apiece before any single orbital is double occupied? As it turns out, the latter situation occurs.

Hund's rule

This states that filled and half-filled shells tend to have additional stability. In some instances, then, for example, the 4s orbitals will be filled before the 3d orbitals.

This rule is applicable only for those elements that have d electrons, and so is less important in organic chemistry (though it is important in organometallic chemistry).

From WP: Hund's rule of maximum multiplicity, often simply referred to as Hund's rule, is a principle of atomic chemistry which states that a greater total spin state usually makes the resulting atom more stable, most commonly manifested in a lower energy state, because it forces the unpaired electrons to reside in different spatial orbitals. A commonly given reason for the increased stability of high multiplicity states is that the different occupied spatial orbitals create a larger average distance between electrons, reducing electron-electron repulsion energy. In reality, it has been shown that the actual reason behind the increased stability is a decrease in the screening of electron-nuclear attractions. Total spin state is calculated as the total number of unpaired electrons + 1, or twice the total spin + 1 written as 2s+1.

Octet rule

The octet rule states that atoms tend to prefer to have eight electrons in their valence shell, so will tend to combine in such a way that each atom can have eight electrons in its valence shell, similar to the electronic configuration of a noble gas. In simple terms, molecules are more stable when the outer shells of their constituent atoms are empty, full, or have eight electrons in the outer shell.

The main exception to the rule is hydrogen, which is at lowest energy when it has two electrons in its valence shell.

Other notable exceptions are aluminum and boron, which can function well with six valence electrons; and some atoms beyond group three on the periodic table that can have over eight electrons, such as sulfur. Additionally, some noble gasses can form compounds when expanding their valence shell.

The other tendency of atoms with regard to their electrons is to maintain a neutral charge. Only the noble gasses have zero charge with filled valence octets. All of the other elements have a charge when they have eight electrons all to themselves. The result of these two guiding principals is the explanation for much of the reactivity and bonding that is observed within atoms; atoms seeking to share electrons in a way that minimizes charge while fulfilling an octet in the valence shell.

Molecular orbitals

In organic chemistry we look at the hybridization of electron orbitals into something called molecular orbitals.

The s and p orbitals in a carbon atom combine into four hybridized orbitals that repel each other in a shape much like that of four balloons tied together. Carbon takes this tetrahedral shape because it only has six electrons which fill the s but only two of the p orbitals.

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When all the s and p orbitals are entirely full the atom's electron clouds form a shape called an octahedral, which is similar to a 3-dimensional diamond in that it is formed by two square pyramids whose bases are placed against each other.

Hybridization

Hybridization refers to the combining of the orbitals of two or more covalently bonded atoms. Depending on how many free electrons a given atom has and how many bonds it is forming, the electrons in the s and the p orbitals will combine in certain manners to form the bonds.

It is easy to determine the hybridization of an atom given a Lewis structure. First, you count the number of pairs of free electrons and the number of sigma bonds (single bonds). Do not count double bonds, since they do not affect the hybridization of the atom. Once the total of these two is determined, the hybridization pattern is as follows:

Sigma Bonds + Electron Pairs     Hybridization
              2                       sp
              3                       sp2
              4                       sp3

The pattern here is the same as that for the electron orbitals, which serves as a memory guide.

Electronegativity

Whenever two atoms form a bond, the nucleus of each atom attracts the other's electrons. Electronegativity is a measure of the strength of this attraction.

Periodic trends

Several traits of atoms are said to have "periodic trends", meaning that different atoms in a period have identifiable relationships to one another based on their position. Is that confusing? Think of the periodic table as a group picture, maybe of a very large basketball team. Each period is a row of players in the picture, and the "photographer" has decided to arrange the "players" by their characteristics. Of course, no conscious effort was made to arrange the periodic table by any characteristic other than number of protons, but some properties are consistent in its layout regardless.

Atomic size is one characteristic that shows a periodic trend. In case of atomic radius the "photographer" (Mendeleev and others since) decided to arrange "players" (atoms) by size with the very shortest and smallest players at the top right. As you go left to right along a row (a period) the atoms get sequentially smaller and smaller. Fluorine is smaller than carbon, and carbon is smaller than magnesium. This is due to the number of protons in the nucleus increasing, while the increasing number of electrons are unable to shield one another from the attractive force of the positive charge from the nucleus.

REMEMBER: largest > Li > Be > B > C > N > O > F > Ne > smallest

Another characteristic with a periodic trend is ionization energy. This is the amount of energy necessary to remove one electron from an atom. Since all the atoms favor an electron configuration of a noble gas, the atoms at the extreme left of the table will give up their first electron most readily. (In almost all cases, a metal will readily give up its first electron.) The halogens, which need only one more electron to fill their outer shells, require a great deal of energy to give up an electron because they would be much more stable if they gained one electron instead. Ionization energy is the opposite of atomic radius, therefore, because it increases from left to right across a period.

REMEMBER: least energy to ionize < Li < Be < B < C < N < O < F < Ne < most energy to ionize

Electronegativity is perhaps the most important periodic trend, and it is not related to ionization energy directly -- but its trend is the same, increasing from left to right. Also, the elements in a group (like the halogen group) gain stability as they grow in atomic number, so the smallest member of an electronegative group is often the most electronegative. In general, it can be said that among periods (rows) or groups (columns) of the periodic table, the closer an element is to fluorine, the more electronegative it will be. For Group VIIA (the aforementioned halogens) of the periodic table, you memorize the following relationships:

REMEMBER: most electronegative > F > Cl > Br > I > least electronegative

And REMEMBER: least electronegative < Li < Be < B < C < N < O < F < most electronegative

(Notice that the noble gas Neon is not on the electronegativity chart. In its non-ionized form, a noble gas is usually treated as if it has no electronegativity at all.)

Electronegativities of atoms common in organic chemistry

• C - 2.5
• H - 2.1
• N - 3.0
• O - 3.5
• P - 2.1
• S - 2.5
• Cl - 3.0
• Br - 2.8
• F - 4.0

Higher numbers represent a stronger attraction of electrons.

When atoms of similar electronegativity bond, a nonpolar covalent bond is the result.

Common nonpolar bonds C-C H-C

When atoms of slightly different electronegativities bond, a polar covalent bond results.

Common polar bonds δ+ C-O δ- δ+ C-N δ- δ- O-H δ+ δ- N-H δ+ δ- and δ+ represent partial charges

When atoms of very different electronegativities bond, an ionic bond results.

Electronegativity content from Wikipedia

Electronegativity is a measure of the ability of an atom or molecule to attract electrons in the context of a chemical bond. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved. Atoms with similar electronegativities will share an electron with each other and form a covalent bond. However, if the difference is too great, the electron will be permanently transferred to one atom and an ionic bond will form. Furthermore, in a covalent bond if one atom pulls slightly harder than the other, a polar covalent bond will form.

The reverse of electronegativity, the ability of an atom to lose electrons, is known as electropositivity.

Two scales of electronegativity are in common use: the Pauling scale (proposed in 1932) and the Mulliken scale (proposed in 1934). Another proposal is the Allred-Rochow scale.

Pauling scale

The Pauling scale was devised in 1932 by Linus Pauling. On this scale, the most electronegative chemical element (fluorine) is given an electronegativity value of 3.98 (textbooks often state this value to be 4.0); the least electronegative element (francium) has a value of 0.7, and the remaining elements have values in between. On the Pauling scale, hydrogen is arbitrarily assigned a value of 2.1 or 2.2.

'δEN' is the difference in electronegativity between two atoms or elements. Bonds between atoms with a large electronegativity difference (greater than or equal to 1.7) are usually considered to be ionic, while values between 1.7 and 0.4 are considered polar covalent. Values below 0.4 are considered non-polar covalent bonds, and electronegativity differences of 0 indicate a completely non-polar covalent bond.

Mulliken scale

The Mulliken scale was proposed by Robert S. Mulliken in 1934. On the Mulliken scale, numbers are obtained by averaging ionization potential and electron affinity. Consequently, the Mulliken electronegativities are expressed directly in energy units, usually electron volts.

Electronegativity trends

Each element has a characteristic electronegativity ranging from 0 to 4 on the Pauling scale. The most strongly electronegative element, fluorine, has an electronegativity of 3.98 while weakly electronegative elements, such as lithium, have values close to 1. The least electronegative element is francium at 0.7. In general, the degree of electronegativity decreases down each group and increases across the periods, as shown below. Across a period, non-metals tend to gain electrons and metals tend to lose them due to the atom striving to achieve a stable octet. Down a group, the nuclear charge has less effect on the outermost shells. Therefore, the most electronegative atoms can be found in the upper, right hand side of the periodic table, and the least electronegative elements can be found at the bottom left. Consequently, in general, atomic radius decreases across the periodic table, but ionization energy increases.

→ Atomic radius decreases → Ionization energy increases → Electronegativity increases →
Group	1	2	3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
Period
1	H 2.20																	He
2	Li 0.98	Be 1.57											B 2.04	C 2.55	N 3.04	O 3.44	F 3.98	Ne
3	Na 0.93	Mg 1.31											Al 1.61	Si 1.90	P 2.19	S 2.58	Cl 3.16	Ar
4	K 0.82	Ca 1.00	Sc 1.36	Ti 1.54	V 1.63	Cr 1.66	Mn 1.55	Fe 1.83	Co 1.88	Ni 1.91	Cu 1.90	Zn 1.65	Ga 1.81	Ge 2.01	As 2.18	Se 2.55	Br 2.96	Kr 3.00
5	Rb 0.82	Sr 0.95	Y 1.22	Zr 1.33	Nb 1.6	Mo 2.16	Tc 1.9	Ru 2.2	Rh 2.28	Pd 2.20	Ag 1.93	Cd 1.69	In 1.78	Sn 1.96	Sb 2.05	Te 2.1	I 2.66	Xe 2.6
6	Cs 0.79	Ba 0.89	*	Hf 1.3	Ta 1.5	W 2.36	Re 1.9	Os 2.2	Ir 2.20	Pt 2.28	Au 2.54	Hg 2.00	Tl 1.62	Pb 2.33	Bi 2.02	Po 2.0	At 2.2	Rn
7	Fr 0.7	Ra 0.9	**	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Uub	Uut	Uuq	Uup	Uuh	Uus	Uuo
Lanthanides	*	La 1.1	Ce 1.12	Pr 1.13	Nd 1.14	Pm 1.13	Sm 1.17	Eu 1.2	Gd 1.2	Tb 1.1	Dy 1.22	Ho 1.23	Er 1.24	Tm 1.25	Yb 1.1	Lu 1.27
Actinides	**	Ac 1.1	Th 1.3	Pa 1.5	U 1.38	Np 1.36	Pu 1.28	Am 1.13	Cm 1.28	Bk 1.3	Cf 1.3	Es 1.3	Fm 1.3	Md 1.3	No 1.3	Lr

Periodic table of electronegativity using the Pauling scale

Bonding

Ionic Bonding

Ionic bonding is when positively and negatively charged ions stick to each other through electrostatic force. These bonds are slightly weaker than covalent bonds and stronger than Van der Waals bonding or hydrogen bonding.

In ionic bonds the electronegativity of the negative ion is so much stronger than the electronegativity of the positive ion that the two ions do not share electrons. Rather, the more electronegative ion assumes full ownership of the electron(s).

Sodium chloride forms crystals with cubic symmetry. In these, the larger chloride ions are arranged in a cubic close-packing, while the smaller sodium ions fill the octahedral gaps between them. Each ion is surrounded by six of the other kind. This same basic structure is found in many other minerals, and is known as the halite structure.

Perhaps the most common example of an ionically bonded substance is NaCl, or table salt. In this, the sodium (Na) atom gives up an electron to the much more electronegative chlorine (Cl) atom, and the two atoms become ions, Na⁺ and Cl^-.The electrostatic bonding force between the two oppositely charged ions extends outside the local area attracting other ions to form giant crystal structures. For this reason most ionically bonded materials are solid at room temperature.

Covalent Bonding

Covalent bonding is close to the heart of organic chemistry. This is where two atoms share electrons in a bond. The goal of each atom is to fill its octet as well as have a formal charge of zero. To do this, atomic nuclei share electrons in the space between them. This sharing also allows the atoms to reach a lower energy state, which stabilizes the molecule. Most reactions in chemistry are due to molecules achieving a lower energy state. Covalent bonds are most frequently seen between atoms with similar electronegativity. In molecules that only have one type of atom, e.g. H₂ or O₂ , the electronegativity of the atoms is necessarily identical, so they cannot form ionic bonds. They always form covalent bonds.

Carbon is especially good at covalent bonding because its electronegativity is intermediate relative to other atoms. That means it can give as well as take electrons as needs warrant.

Covalently bonded compounds have strong internal bonds but weak attractive forces between molecules. Because of these weak attractive forces, the melting and boiling points of these compounds are much lower than compounds with ionic bonds. Therefore, such compounds are much more likely to be liquids or gases at room temperature than ionically bonded compounds.

In molecules formed from two atoms of the same element, there is no difference in the electronegativity of the bonded atoms, so the electrons in the covalent bond are shared equally, resulting in a completely non-polar covalent bond. In covalent bonds where the bonded atoms are different elements, there is a difference in electronegativities between the two atoms. The atom that is more electronegative will attract the bonding electrons more toward itself than the less electronegative atom. The difference in charge on the two atoms because of the electrons causes the covalent bond to be polar. Greater differences in electronegativity result in more polar bonds. Depending on the difference in electronegativities, the polarity of a bond can range from non-polar covalent to ionic with varying degrees of polar covalent in between. An overall imbalance in charge from one side of a molecule to the other side is called a dipole moment. Such molecules are said to be polar. For a completely symmetrical covalently bonded molecule, the overall dipole moment of the molecule is zero. Molecules with larger dipole moments are more polar. The most common polar molecule is water.

Bond Polarity and Dipole Moment

The ideas of bond polarity and dipole moment play important roles in organic chemistry.

If you look at the image of methane on the right, the single most important aspect of it in terms of bond polarity is that it is a symmetric molecule. It has 4 hydrogens, all bonded at 109.5° from the other, and all with precisely the same bond angle. Each carbon-hydrogen bond is slightly polar (hydrogen has an electronegativity of 2.1, carbon 2.5), but because of this symmetry, the polarities cancel each other out and overall, methane is a non-polar molecule.

The distinction is between Bond Polarity and Molecular polarity. The total polarity of a molecule is measured as Dipole Moment. The actual calculation of dipole moment isn't really necessary so much as an understanding of what it means. Frequently, a guesstimate of dipole moment is pretty easy once you understand the concept and until you get into the more advanced organic chemistry, exact values are of little value.

Basically, the molecular polarity is, essentially, the summation of the vectors of all of the bond polarities in a molecule.

Van der Waals Bonding

Van der Waals bonding is the collective name for three types of interactions:

1. Permanent Dipole interactions: these are the electrostatic attractive forces between two dipoles, these are responsible for fluromethane's (CH3F) high boiling point (about -15 deg C) compared to Nitrogen (about -180 deg C).
2. Permanent dipole / induced dipole: these are the interactions between a permanent dipole and another molecule, causing the latter molecule's electron cloud to be distorted and thus have an induced dipole itself. These are much weaker than the permanent dipole / dipole interactions. These forces occur in permanent dipole-molecules, and in mixtures of permanent dipole and dipole free molecules.
3. Instantaneous dipole / induced dipole: At any specific moment the electron cloud is not necesarily symetrical, this instantaneous dipole then induces a dipole in another molecule and they are attracted, this is the weakest of all molecular interactions.

A Dipole is caused by an atom or molecule fragment having a higher electronegativity (this is a measure of its effective nuclear charge, and thus the attraction of the nucleus by electrons) than one to which it is attached. This means that it pulls electrons closer to it, and has a higher share of the electrons in the bond. Dipoles can cancel out by symmetry, eg: Carbon dioxide (O=C=O) is linear so there is no dipole, but the charge distribution is asymmetric causing a quadripole moment (this acts similarly to a dipole, but is much weaker).

Organometallic Compounds and Bonding

Organometallic chemistry combines aspects of inorganic chemistry and organic chemistry, because organometallic compounds are chemical compounds containing bonds between carbon and a metal or metalloid element. Organometallic bonds are different from other bonds in that they are not either truly covalent or truly ionic, but each type of metal has individual bond character. Cuprate (copper) compounds, for example, behave quite differently than Grignard reagents (magnesium), and so beginning organic chemists should concentrate on how to use the most basic compounds mechanistically, while leaving the explanation of exactly what occurs at the molecular level until later and more in-depth studies in the subject.

Basic organometallic interactions are discussed fully in a later chapter.

Electron dot structures & formal charge

Electron Dot Structures

Electron dot structures, also called Lewis structures, give a representation of the valence electrons surrounding an atom.

Each valence electron is represented by one dot, thus, a lone atom of hydrogen would be drawn as an H with one dot, whereas a lone atom of Helium would be drawn as an He with two dots, and so forth.

Representing two atoms joined by a covalent bond is done by drawing the atomic symbols near to each other, and drawing a single line to represent a shared pair of electrons. It is important to note: a single valence electron is represented by a dot, whereas a pair of electrons is represented by a line.

The covalent compound hydrogen fluoride, for example, would be represented by the symbol H joined to the symbol F by a single line, with three pairs (six more dots) surrounding the symbol F. The line represents the two electrons shared by both hydrogen and fluorine, whereas the six paired dots represent fluorine's remaining six valence electrons.

Dot structures are useful in illustrating simple covalent molecules, but the limitations of dot structures become obvious when diagramming even relatively simple organic molecules. The dot structures have no ability to represent the actual physical orientation of molecules, and they become overly cumbersome when more than three or four atoms are represented.

Lewis dot structures are useful for introducing the idea of covalence and bonding in small molecules, but other model types have much more capability to communicate chemistry concepts.

Drawing electron dot structures

Some examples of electron dot structures for a few commonly encountered molecules from inorganic chemistry.

A note about Gilbert N. Lewis

Lewis was born in Weymouth, Massachusetts as the son of a Dartmouth-graduated lawyer/broker. He attended the University of Nebraska at age 14, then three years later transferred to Harvard. After showing an initial interest in Economics, Gilbert Newton Lewis earned first a B.A. in Chemistry, and then a Ph.D. in Chemistry in 1899.

For a few years after obtaining his doctorate, Lewis worked and studied both in the United States and abroad (including Germany and the Phillipines) and he was even a professor at M.I.T. from 1907 until 1911. He then went on to U.C. Berkeley in order to be Dean of the College of Chemistry in 1912.

In 1916 Dr. Lewis formulated the idea that a covalent bond consisted of a shared pair of electrons. His ideas on chemical bonding were expanded upon by Irving Langmuir and became the inspiration for the studies on the nature of the chemical bond by Linus Pauling.

In 1923, he formulated the electron-pair theory of acid-base reactions. In the so-called Lewis theory of acids and bases, a "Lewis acid" is an electron-pair acceptor and a "Lewis base" is an electron-pair donor.

In 1926, he coined the term "photon" for the smallest unit of radiant energy.

Lewis was also the first to produce a pure sample of deuterium oxide (heavy water) in 1933. By accelerating deuterons (deuterium nuclei) in Ernest O. Lawrence's cyclotron, he was able to study many of the properties of atomic nuclei.

During his career he published on many other subjects, and he died at age 70 of a heart attack while working in his laboratory in Berkeley. He had one daughter and two sons; both of his sons became chemistry professors themselves.

Formal Charge

The formal charge of an atom is the charge that it would have if every bond were 100% covalent (non-polar). Formal charges are computed by using a set of rules and are useful for accounting for the electrons when writing a reaction mechanism, but they don't have any intrinsic physical meaning. They may also be used for qualitative comparisons between different resonance structures (see below) of the same molecule, and often have the same sign as the partial charge of the atom, but there are exceptions.

The formal charge of an atom is computed as the difference between the number of valence electrons that a neutral atom would have and the number of electrons that "belong" to it in the Lewis structure when one counts lone pair electrons as belonging fully to the atom, while electrons in covalent bonds are split equally between the atoms involved in the bond. The total of the formal charges on an ion should be equal to the charge on the ion, and the total of the formal charges on a neutral molecule should be equal to zero.

For example, in the hydronium ion, H₃O⁺, the oxygen atom has 5 electrons for the purpose of computing the formal charge—2 from one lone pair, and 3 from the covalent bonds with the hydrogen atoms. The other 3 electrons in the covalent bonds are counted as belonging to the hydrogen atoms (one each). A neutral oxygen atom has 6 valence electrons (due to its position in group 16 of the periodic table); therefore the formal charge on the oxygen atom is 6 – 5 = +1. A neutral hydrogen atom has one electron. Since each of the hydrogen atoms in the hydronium atom has one electron from a covalent bond, the formal charge on the hydrogen atoms is zero. The sum of the formal charges is +1, which matches the total charge of the ion.

Formal Charge: number of valence electrons for an atom - (number of lone pair electrons + number electrons in bonds/2)

In chemistry, a formal charge (FC) on an atom in a molecule is defined as:

FC = number of valence electrons of the atom - ( number of lone pair electrons on this atom + total number of electrons participating in covalent bonds with this atom / 2).

When determining the correct Lewis structure (or predominant resonance structure) for a molecule, the structure is chosen such that the formal charge on each of the atoms is minimized.

Examples

carbon in methane 

Nitrogen in  

double bonded oxygen in  

single bonded oxygen in  

Methane (CH4): black is carbon, white is hydrogen

Nitrogen dioxide (NO2): blue is nitrogen, red is oxygen

Resonance

Resonance

Resonance refers to structures that are not easily represented by a single electron dot structure but that are intermediates between two or more drawn structures.

Resonance is easily misunderstood in part because of the way certain chemistry textbooks attempt to explain the concept. In science, analogies can provide an aid to understanding, but analogies should not be taken too literally. It is sometimes best to use analogies to introduce a topic, but then explain the differences and inevitable complications as further details on a complicated subject. This is the case for resonance.

Just as entropic principles cannot be applied to individual molecules, so it is impossible to say whether or not any given individual molecule with a resonance structure is literally in one configuration or another. The actual situation on the molecular scale is that each configuration of the molecule contributes a percentage to the possible configurations, resulting in a "blend" of the possible structures. Changes in molecular shape occur so rapidly, and on such a tiny scale, that the actual physical locations of individual electrons cannot be precisely known (due to Heisenberg's Uncertainty Principle). The result of all that complexity is simply this: molecules with resonance structures are treated as mixtures of their multiple forms, with a greater percentage of probability given to the most stable configurations.

The nuclei of the atoms are not moving when they are represented by resonance structure drawings. Rather, the electrons are portrayed as if they were moving instead. The true situation is that no one can say for certain exactly where any individual electron is at any specific moment, but rather electron location can be expressed as a probability only. What a dot structure is actually showing is where electrons almost certainly are located, therefore resonance structures indicate a split in those same probabilities. Chemists are absolutely certain where electrons are located when one carbon bonds four hydrogens (methane), but it is less certain where precisely any given electron is located when six carbons bond six hydrogens in a ring structrue (benzene). Resonance is an expression of this uncertainty, and is therefore the average of probable locations.

Resonance structures are stabilizing in molecules because they allow electrons to lengthen their wavelengths and thereby lower their energy. This is the reason that benzene (C₆H₆) has a lower heat of formation than organic chemists would predict, not accounting for resonance. Other aromatic molecules have a similar stability, which leads to an overall entropic preference for aromaticity (a subject that will be covered fully in a later chapter). Resonance stability plays a major role in organic chemistry due to resonant molecules' lower energy of formation, so students of organic chemistry should understand this effect and practice spotting molecules stabilized by resonant forms.

In the Lewis structures above, carbonate (CO₃) has a resonance structure. Using laboratory procedures to measure the bond length of each bond, we do not find that one bond is shorter than the two others (remember, double bonds are shorter than single bonds), but instead that all bonds are of the same length somewhere between the length of typical double and single bonds.

Resonance Structures

Resonance structures are diagrammatic tools used predominately in organic chemistry to symbolize resonant bonds between atoms in molecules. The electron density of these bonds is spread over the molecule, also known as the delocalization of electrons. Resonance contributors for the same molecule all have the same chemical formula and same sigma framework, but the pi electrons will be distributed differently among the atoms. Because Lewis dot diagrams often cannot represent the true electronic structure of a molecule, resonance structures are often employed to approximate the true electronic structure. Resonance structures of the same molecule are connected with a double-headed arrow. While organic chemists use resonance structures frequently, they are used in inorganic structures, with nitrate as an example.

Key characteristics

The key elements of resonance are:

• Resonance occurs because of the overlap of orbitals. Double bonds are made up of pi bonds, formed from the overlap of 2p orbitals. The electrons in these pi orbitals will be spread over more than two atoms, and hence are delocalized.
• Both paired and unshared electrons may be delocalized, but all the electrons must be conjugated in a pi system.
• If the orbitals do not overlap (such as in orthogonal orbitals) the structures are not true resonance structures and do not mix.
• Molecules or species with resonance structures are generally considered to be more stable than those without them. The delocalization of the electrons lowers the orbital energies, imparting this stability. The resonance in benzene gives rise to the property of aromaticity. The gain in stability is called the resonance energy.
• All resonance structures for the same molecule must have the same sigma framework (sigma bonds form from the "head on" overlap of hybridized orbitals). Furthermore, they must be correct Lewis structures with the same number of electrons (and consequent charge) as well as the same number of unpaired electrons. Resonance structures with arbitrary separation of charge are unimportant, as are those with fewer covalent bonds. These unimportant resonance structures only contribute minimally (or not at all) to the overall bonding description; however, they are important in some cases such as for a carbonyl group.
• The hybrid structure is defined as the superposition of the resonance structures. A benzene ring is often shown with a circle inside a hexagon (in American texts) rather than alternating double bonds — the latter example misrepresents the electronic structure. Bonds with broken bond orders are often displayed as double bonds with one solid and one dashed line.

What resonance is not

Significantly, resonance structures do not represent different, isolatable structures or compounds. In the case of benzene, for example, there are two important resonance structures - which can be thought of as cyclohexa-1,3,5-trienes. There are other resonance forms possible, but because they are higher in energy than the triene structures (due to charge separation or other effects) they are less important and contribute less to the "real" electronic structure (average hybrid). However, this does not mean there are two different, interconvertable forms of benzene; rather, the true electronic structure of benzene is an average of the two structures. The six carbon-carbon bond lengths are identical when measured, which would be invalid for the cyclic triene. Resonance should also not be confused with a chemical equilibrium or tautomerism which are equilibria between compounds that have different sigma bonding patterns. Hyperconjugation is a special case of resonance.

History

The concept of resonance was introduced by Linus Pauling in 1928. He was inspired by the quantum mechanical treatment of the H₂⁺ ion in which an electron is located between two hydrogen nuclei. The alternative term mesomerism popular in German and French publications with the same meaning was introduced by Christopher Ingold in 1938 but did not catch on in the English literature. The current concept of Mesomeric effect has taken on a related but different meaning. The double headed arrow was introduced by the German chemist Arndt (also responsible for the Arndt-Eistert synthesis) who preferred the German phrase zwischenstufe or intermediate phase.

Due to confusion with the physical meaning of the word resonance, as no elements do actually appear to be resonating, it is suggested to abandon the term resonance in favor of delocalization ^[1]. Resonance energy would become delocalization energy and a resonance structure becomes contributing structure. The double headed arrows would get replaced by commas.

Examples

The ozone molecule is represented by two resonance structures in the top of scheme 2. In reality the two terminal oxygen atoms are equivalent and the hybrid structure is drawn on the right with a charge of -1/2 on both oxygen atoms and partial double bonds. The concept of benzene as a hybrid of two conventional structures (middle scheme 2) was a major breakthrough in chemistry made by Kekule, and the two forms of the ring which together represent the total resonance of the system are called Kekule structures. In the hybrid structure on the right the circle replaces three double bonds. The allyl cation (bottom scheme 2) has two resonance forms and in the hybrid structure the positive charge is delocalized over the terminal methylene groups.

See also

• Delocalization

References

1. ^  If It's Resonance, What Is Resonating? Kerber, Robert C. . J. Chem. Educ. 2006 83 223. Abstract
2. (Much of this text originally from http://en.wikipedia.org/w/index.php?title=Resonance_%28chemistry%29&oldid=41962377

Acids and bases

Arrhenius Definition: Hydroxide and Hydronium Ions

The first and earliest definition of acids and bases was proposed in the 1800's by Swedish scientist Svante Arrhenius, who said that an acid was anything that dissolved in water to yield H⁺ ions (like stomach acid HCl, hydrochloric acid), and a base was anything that dissolved in water to give up OH^- ions (like soda lye NaOH, sodium hydroxide). Acids and bases were already widely used in various occupations and activities of the time, so Arrhenius' definition merely attempted to explained well-known and long-observed phenomenon.

Although simple, at the time this definition of the two types of substances was significant. It allowed chemists to explain certain reactions as ion chemistry, and it also expanded the ability of scientists of the time to predict certain chemical reactions. The definition left a great deal wanting, however, in that many types of reactions that did not involve hydroxide or hydronium ions directly remained unexplained.

Many general chemistry classes (especially in the lower grades or introductory levels) still use this simple definition of acids and bases today, but modern organic chemists make further distinctions between acids and bases than the distinctions provided under Arrhenius's definition.

Brønsted-Lowry Acids and Bases: Proton donors and acceptors

A new definition for acids and bases, building upon the one already proposed by Arrhenius, was brought forth independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. The new definition did not depend on a substance's dissolution in water for definition, but instead suggested that a substance was acidic if it readily donated a proton (H⁺) to a reaction and a substance was basic if it accepted a proton in a reaction.

Definiton of Brønsted-Lowry Acid and Base An acid is any proton donor and a base is any proton acceptor.

The major advantage of the updated definition was that it was not limited to aqueous solut