Definitions

Many of the reactions that occur in organic chemistry are either acid–base reactions themselves, or they involve an acid–base reaction at some stage. In general, both acids and bases are electrolytes, and they dissolve in water to give solutions that conduct electricity.

An electrolyte is a substance, which dissociates in a solvent to form an ionic solution. Strong electrolytes dissociate almost completely. Weak electrolytes form an equilibrium between the associated and dissociated forms and are only partly dissociated.

The Bronsted–Lowry Definition of Acids and Bases

In 1923, both T.M. Lowry and J.N. Bronsted proposed that

"An acid can donate a proton to another substance, and a base can accept a proton from another substance."

The molecule or ion that forms when an acid loses its proton is called the conjugate base of that acid. The chloride ion, therefore, is the conjugate base of HCl. The molecule or ion that forms when a base accepts a proton is called the conjugate base of that base. The NH₄⁺ ion, therefore, is the conjugate acid of Ammonia.

Let us consider, as an example of this concept, the reaction that occurs when gaseous hydrogen chloride dissolves in water: Hydrogen chloride, a very strong acid, transfers its proton to water. Water acts as a base and accepts the proton. The products that result from this reaction are a hydronium ion (H₃O⁺) and a chloride ion (Cl⁻). Other strong acids that completely transfer a proton when dissolved in water are hydrogen iodide, hydrogen bromide, and sulfuric acid.

Because sulfuric acid has two protons that it can transfer to a base, it is called a diprotic (or dibasic) acid.

Lewis Definition of Acids and Bases

In the same year as Bronsted and Lowry, Lewis proposed an even more general definition based on the sharing of electron pairs: an acid is able to accept a pair of electrons to form a new bond; a base is able to donate a pair of electrons to form a new bond. Aluminum chloride, for example reacts with ammonia in the same way that a proton does.

A Lewis acid needs an empty valence orbital for the electron pair from the base to go into. Metal cations like Cu²⁺ attract electrons because of their positive charge and have at least one empty orbital. Molecules with an unfilled octet, like BF₃, are strong Lewis acids. Some neutral molecules such as the oxides of non–metals are also Lewis acids. One example of this is carbon dioxide.

Because oxygen is so electronegative, the electrons in the double bonds are drawn away from carbon, giving carbon a slightly positive charge. Carbon then becomes a good site to attract electron pairs.

A Lewis base has at least one unshared pair of valence electrons as does a Bronsted–Lowry base (otherwise it wouldn't be able to accept a proton). Some examples of Lewis bases (which are also Bronsted–Lowry bases):

Let us consider the following reactions. The donation of the electron pair of ammonia (the Lewis base).

In this example, aluminum chloride accepts the electron pair of ammonia just as a proton does, by using it to form a covalent bond to the nitrogen atom. This happens because the central aluminum atom has only a sextet of electrons and is, therefore, electron deficient. When it accepts the electron pair, aluminum chloride is, in the Lewis definition, acting as an acid.

Bases are much the same in the Lewis theory and the Bronsted–Lowry theory, because in the Bronsted–Lowry theory a base must donate a pair of electrons in order to accept a proton. The Lewis theory, by virtue of its broader definition of acids, allows acid–base theory to include all of the Bronsted – Lowry reactions and a great many others.

Any electron–deficient atom can act as a Lewis acid. Many compounds containing Group 3A elements such as boron and aluminum are Lewis acids because Group 3A atoms have only a sextet of electrons in their outer shell. Many other compounds that have atoms with vacant orbitals also act as Lewis acids. Zinc and iron(III) halides (ferric halides) are frequently used as Lewis acids in organic reactions. Two examples that we shall study later are the following:

Amphiprotism Definition

A prophylactic (also called amphoteric) substance is capable of donating or accepting protons, depending on the reaction conditions. Water is the most common amphoteric species, but not the only one. Liquid ammonia (NH₃) is also amphiprotic, as are the anions of polyprotic acids. Consider how a substance can both accept and donate an electron pair. A polyprotic acid is one, which has more than one proton to donate (eg: H₂SO₄). It is an amphiprotic substance as well.

Strength of Acids and Bases

k_a and pk_a:

In contrast to the strong acids, such as HCl and H₂SO₄, acetic acid is a much weaker acid. When acetic acid dissolves in water, the following reaction does not proceed to completion.

Experiments show that in a 0.1 Msolution of acetic acid at 25°C only about 1 % of the acetic acid molecules ionize by transferring their protons to water.

The Acidity Constant, K_a Because the reaction that occurs in an aqueous solution of acetic acid is an equilibrium, we can describe it with an expression for the equilibrium constant.

For dilute aqueous solutions, the concentration of water is essentially constant (∼ 55.5M), so we can rewrite the expression for the equilibrium constant in terms of a new constant (K_a)called the acidity constant.

At 25°C, the acidity constant for acetic acid is 1.76 × 10⁻⁵.

We can write similar expressions for any weak acid dissolved in water. Using a generalized hypothetical acid (HA) the reaction in water is

and the expression for the acidity constant is

Because the concentrations of the products of the reaction are written in the numerator and the concentration of the undissociated acid in the denominator, a large value of K_ameans the acid is a strong acid, and a small value of K_a means the acid is a weak acid. If the K_a is greater than 10, the acid will be, for all practical purposes, completely dissociated in water. K_a is an indicator of acid strength.

Acidity Negative Logarithm

Acidity PK_a

Chemists usually express the acidity constant, K_a, as its negative logarithm, pK_a.

This is analogous to expressing the hydronium ion concentration as pH.

For acetic acid the pK_a is 4.75

Notice that there is an inverse relationship between the magnitude of the pK_a and the strength of the acid. The larger the value of the pK_a, the weaker is the acid. For example, acetic acid with a pK_a = 4.75 is a weaker acid than trifluoroacetic acid with a pK_a = 0 (K_a = 1). Hydrochloric acid with a pK_a = − 7 (K_a = 10⁷) is a far stronger acid than trifluoroacetic acid. (It is understood that a positive pK_a is larger than a negative pK_a.)

Below table lists pK_a values for a selection of acids relative to water as the base. The values in the middle pK_a range of the table are the most accurate because they can be measured in aqueous solution. Special methods must be used to estimate the pK_a values for the very strong acids at the top of the table and for the very weak acids at the bottom. *The pK_a values for these very strong and weak acids are therefore approximate. All of the acids that we shall consider will have strengths in between that of ethane (an extremely weak acid) and that of HSbF₆ (an acid that is so strong that it is called a "superacid").

Relative strengths of Selected Acids and Their Conjugate Bases.

Water, itself, is a very weak acid and undergoes self– ionization even in the absence of acids and bases.

In pure water at 25°C, the concentrations of hydronium and hydroxide ions are equal to 10⁻⁷M. Since the concentration of water in pure water is 55.5M, we can calculate the K_a for water.

Predicting the Strength of Bases

There is a principle that allows us to estimate the strengths of bases. Simply stated, the principle is this:The stronger the acid, the weaker will be its conjugate base.

We can, therefore, relate the strength of a base to the pK_aof its conjugate acid. The larger the pK_a of the conjugate acid, the stronger is the base. Consider the following as examples:

Hydroxide ion is the strongest base of these three bases because its conjugate acid, water, is the weakest acid. Amines are like ammonia in that they are weak bases. Dissolving ammonia in water brings about the following equilibrium.

Amines are like ammonia in that they are weak bases

Dissolving methylamine in water causes the establishment of a similar equilibrium. Again we can relate the basicity of these substances to the strength of their conjugate acids. The conjugate acid of ammonia is the ammonium ion, NH₄⁺. The pK_a of the ammonium ion is 9.2. The conjugate acid of methylamine is the CH₃NH₃⁺ ion. This ion, called the methylaminum ion, has a pK_a = 10.6. Since the conjugate acid of methylamine is a weaker acid than the conjugate acid of ammonia, we can conclude that methylamine is a stronger base than ammonia.

Trends in acidity

The strength of an acid depends on the extent to which a proton can be separated from it and transferred to a base. Removing the proton involves breaking a bond to the proton, and it involves making the conjugate base more electrically negative. When we compare compounds in a vertical column of the, periodic table, the strength of the bond to the proton is the dominating effect. The acidities of the hydrogen halides furnish an example:

Acidity increases as we descend a vertical column: H–F is the weakest acid and H–I is the strongest. The important factor is the strength of the H–X bond; the stronger the bond, the weaker is the acid. The H–F bond is by far the strongest and the H–I bond is the weakest. Because HI, HBr, and HCl are such strong acids, their conjugate bases (I ⁻, Br ⁻, Cl ⁻) are all very weak bases. The fluoride ion is considerably more basic. Overall the basicity of the halide ions increases in the following way:

We see the same trend of acidities and basicities in other vertical columns of the periodic table. Consider, for example, the column headed by oxygen:

Here the strongest bond is the O–H bond and H₂O is the weakest acid; the weakest bond is the Se–H bond and H₂Se is the strongest acid.

Acidity increases from left to right when we compare compounds in the same horizontal row of the periodic table.Bond strengths are roughly the, same, and the dominant factor becomes the electro negativity of the atom bonded to the hydrogen. The electronegativity of this atom affects acidity in two related ways. It affects the polarity of the bond to the proton and it affects the relative stability of the anion (conjugate base) that forms when the proton is lost. Let us compare two hypothetical acids, H–A and H–B.

Let us assume that A is more electronegative than B. The greater electronegativity of A will cause atom A to be more negative than atom B, and the hydrogen (proton) of H–A will be more positive than that of H–B. The proton of H–A, consequently, will be held less strongly, and it will separate and be transferred to a base more readily. The greater electronegativity of A will also mean that atom A will acquire a negative charge more readily than B, and that the A⁻ anion will be more stable than the B⁻ anion. H–A, therefore, will be the stronger acid.

We can see an example of this effect when we compare the acidities of the compounds CH₄, NH₃, H₂O, and HF. These compounds are all hydrides of first–row elements, and electronegativity increases across a row of the periodic table from left to right.

Because fluorine is the most electronegative, the bond in H–F is most polarized, and the proton in H–F is the most positive. Therefore, H–F loses a proton most readily and is the most acidic:

Almost no positive charge is evident at the hydrogens of methane. Very little positive charge is present at the hydrogens of ammonia. This is consistent with the weak electronegativity of both carbon and nitrogen, and hence with the behavior of methane and ammonia as exceedingly weak acids (pK_as of 48 and 38 respectively). Water shows significant positive charge at its hydrogens (pK_a more than 20 units lower than ammonia), and hydrogen fluoride clearly has the highest amount of positive charge at its hydrogen (pK_a of 3.2), resulting in strong acidity.

Because H–F is the strongest acid, its conjugate base, the fluoride ion (F⁻), will be the weakest base. Fluorine is the most electronegative atom and it accommodates the negative charge most readily.

The Methanide Ion Is The Least Stable Anion

The methanide ion (CH₃ ⁻) is the least stable anion of the four, because carbon being the least electronegative element is least able to accept the negative charge. The methanide ion, therefore, is the strongest base. The methanide ion, a carbanion, and the amide ion (NH₂ ⁻) are exceedingly strong bases because they are the conjugate bases of extremely weak acids.

Energy changes

Chemical energy is a form of potential energy. It exists because attractive and repulsive electrical forces exist between different pieces of the molecules. Nuclei attract electrons, nuclei repel each other, and electrons repel each other.

It is usually impractical (and often impossible) to describe the absolute amount of potential energy contained by a substance. Thus we usually think in terms of its relative potential energy. We say that one system has more or less potential energy than another. Another term that chemists frequently use in this context is ‘stability’ or relative stability. The relative stability of a system is inversely related to its relative potential energy. The more potential energy an object has, the less stable it is. Consider, as an example, the relative potential energy and the relative stability of snow when it lies high on a mountainside and when it lies serenely in the valley below. Because of the attractive force of gravity, the snow high on the mountain has greater potential energy and is much less stable than the snow in the valley. This greater potential energy of the snow on the mountainside can become converted into the enormous kinetic energy of a landslide. By contrast, the snow in the valley, with its lower potential energy and with its greater stability, is incapable of releasing such energy.

Potential Energy and Covalent Bonds

Atoms and molecules possess potential energy – often called chemical energy, that can be released as heat when they react. Because heat is associated with molecular motion, this release of heat results from a change from potential energy to kinetic energy.

From the standpoint of covalent bonds, the state of greatest potential energy is the state of free atoms, the state in which the atoms are not bonded to each other at all. This is true because the formation of a chemical bond is always accompanied by the lowering of the potential energy of the atoms. Consider as an example the formation of hydrogen molecules from hydrogen atoms:

Formation Of Hydrogen Molecules From Hydrogen Atoms

The potential energy of the atoms decreases by 435 kJ mol⁻¹ as the covalent bond forms. This potential energy change is illustrated graphically in the figure below.

A convenient way to represent the relative potential energies of molecules is in terms of their relative enthalpies or heat contents, H. (Enthalpy comes from en + Gk: thalpein to heat.) The difference in relative enthalpies of reactants and products in a chemical change is called the enthalpy change and is symbolized by ΔH°. [The δ (delta) in front of a quantity usually means the difference, or change, in the quantity. The superscript ° indicates that the measurement is made under standard conditions.] By convention, the sign of ΔH° for exothermic reactions (those evolving heat) is negative. Endothermic reactions (those that absorb heat) have a positive ΔH°. The heat of reaction, ΔH°, measures the change in enthalpy of the atoms of the reactants as they are converted to products. For an exothermic reaction the atoms have a smaller enthalpy as products than they do as reactants. For endothermic reactions, the reverse is true.