Acids and Bases //د// عاطف خليفه
 

The Brnsted, or Brnsted-Lowry, model is based on a simple assumption: Acids donate H+ ions to another ion or molecule, which acts as a base. The dissociation of water, for example, involves the transfer of an H+ ion from one water molecule to another to form H3O+ and OH- ions.

According to this model, HCl doesn't dissociate in water to form H+ and Cl+ ions. Instead, an H+ ion is transferred from HCl to a water molecule to form H3O+ and Cl- ions, as shown in the figure below.
Because it is a proton, an H+ ion is several orders of magnitude smaller than the smallest atom. As a result, the charge on an isolated H+ ion is distributed over such a small amount of space that this H+ ion is attracted toward any source of negative charge that exists in the solution. Thus, the instant that an H+ ion is created in an aqueous solution, it bonds to a water molecule. The Brnsted model, in which H+ ions are transferred from one ion or molecule to another, therefore makes more sense than the Arrhenius theory, which assumes that H+ ions exist in aqueous solution.
Even the Brnsted model is naive. Each H+ ion that an acid donates to water is actually bound to four neighboring water molecules, as shown in the figure below.
A more realistic formula for the substance produced when an acid loses an H+ ion is therefore H(H2O)4+, or H9O4+. For all practical purposes, however, this substance can be represented as the H3O+ ion.
The reaction between HCl and water provides the basis for understanding the definitions of a Brnsted acid and a Brnsted base. According to this theory, an H+ ion is transferred from an HCl molecule to a water molecule when HCl dissociates in water.
HCl acts as an H+-ion donor in this reaction, and H2O acts as an H+ ion-acceptor. A Brnsted acid is therefore any substance (such as HCl) that can donate an H+ ion to a base. A Brnsted base is any substance (such as H2O) that can accept an H+ ion from an acid.
There are two ways of naming the H+ ion. Some chemists call it a hydrogen ion; others call it a proton. As a result, Brnsted acids are known as either hydrogen-ion donors or proton donors. Brnsted bases are hydrogen-ion acceptors or proton acceptors.
From the perspective of the Brnsted model, reactions between acids and bases always involve the transfer of an H+ ion from a proton donor to a proton acceptor. Acids can be neutral molecules.

They can also be positive ions

or negative ions.

The Brnsted theory therefore expands the number of potential acids. It also allows us to decide which compounds are acids from their chemical formulas. Any compound that contains hydrogen with an oxidation number of +1 can be an acid. Brnsted acids include HCl, H2S, H2CO3, H2PtF6, NH4+, HSO4-, and HMnO4.
Brnsted bases can be identified from their Lewis structures. According to the Brnsted model, a base is any ion or molecule that can accept a proton. To understand the implications of this definition, look at how the prototypical base, the OH- ion, accepts a proton.
The only way to accept an H+ ion is to form a covalent bond to it. In order to form a covalent bond to an H+ ion that has no valence electrons, the base must provide both of the electrons needed to form the bond. Thus, only compounds that have pairs of nonbonding valence electrons can act as H+-ion acceptors, or Brnsted bases.
The following compounds, for example, can all act as Brnsted bases because they all contain nonbonding pairs of electrons.
The Brnsted model expands the list of potential bases to include any ion or molecule that contains one or more pairs of nonbonding valence electrons. The Brnsted definition of a base applies to so many ions and molecules that it is almost easier to count substances, such as the following, that can't be Brnsted bases because they don't have pairs of nonbonding valence electrons.
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The Role of Water in the Brnsted Theory

The Brnsted theory explains water's role in acid-base reactions.

• Water dissociates to form ions by transferring an H+ ion from one molecule acting as an acid to another molecule acting as a base.

H2O(l)+H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gifH3O+(aq)+ OH-(aq)acid base

• Acids react with water by donating an H+ ion to a neutral water molecule to form the H3O+ ion.

HCl(g)+H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gifH3O+(aq)+ Cl-(aq) acid base

• Bases react with water by accepting an H+ ion from a water molecule to form the OH- ion.

NH3(aq)+H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gifNH4+(aq)+ OH-(aq)base acid

• Water molecules can act as intermediates in acid-base reactions by gaining H+ ions from the acid

HCl(g) +H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gifH3O+(aq)+ Cl-(aq)

and then losing these H+ ions to the base.

NH3(aq)+H3O+(aq)http://chemed.chem.purdue.edu/genche...s/equilibr.gifNH4+(aq)+ H2O(l)

The Brnsted model can be extended to acid-base reactions in other solvents. For example, there is a small tendency in liquid ammonia for an H+ ion to be transferred from one NH3 molecule to another to form the NH4+ and NH2- ions.

2 NH3http://chemed.chem.purdue.edu/genche...s/equilibr.gifNH4++ NH2-

By analogy to the chemistry of aqueous solutions, we conclude that acids in liquid ammonia include any source of the NH4+ ion and that bases include any source of the NH2- ion.
The Brnsted model can even be extended to reactions that don't occur in solution. A classic example of a gas-phase acid-base reaction is encountered when open containers of concentrated hydrochloric acid and aqueous ammonia are held next to each other. A white cloud of ammonium chloride soon forms as the HCl gas that escapes from one solution reacts with the NH3 gas from the other.

HCl(g)+ NH3(g)http://chemed.chem.purdue.edu/genche...s/equilibr.gifNH4Cl(s)

This reaction involves the transfer of an H+ ion from HCl to NH3 and is therefore a Brnsted acid-base reaction, even though it occurs in the gas phase

Acids and bases exist as conjugate acid-base pairs. The term conjugate comes from the Latin stems meaning "joined together" and refers to things that are joined, particularly in pairs, such as Brnsted acids and bases.
Every time a Brnsted acid acts as an H+-ion donor, it forms a conjugate base. Imagine a generic acid, HA. When this acid donates an H+ ion to water, one product of the reaction is the A- ion, which is a hydrogen-ion acceptor, or Brnsted base.

Conversely, every time a base gains an H+ ion, the product is a Brnsted acid, HA.

Acids and bases in the Brnsted model therefore exist as conjugate pairs whose formulas are related by the gain or loss of a hydrogen ion.
Our use of the symbols HA and A- for a conjugate acid-base pair does not mean that all acids are neutral molecules or that all bases are negative ions. It signifies only that the acid contains an H+ ion that isn't present in the conjugate base. Brnsted acids or bases can be neutral molecules, positive ions, or negative ions. Various Brnsted acids and their conjugate bases are given in the table below.

A compound can be both a Brnsted acid and a Brnsted base. H2O, OH-, HSO4-, and NH3, for example, can be found in both columns in the table above. Water is the perfect example of this behavior because it simultaneously acts as an acid and a base when it forms the H3O+ and OH- ions.

Many hardware stores sell "muriatic acid" http://chemed.chem.purdue.edu/genche...raphics/em.gif a 6 M solution of hydrochloric acid HCl(aq)http://chemed.chem.purdue.edu/genche...raphics/em.gifto clean bricks and concrete. Grocery stores sell vinegar, which is a 1 M solution of acetic acid: CH3CO2H. Although both substances are acids, you wouldn't use muriatic acid in salad dressing, and vinegar is ineffective in cleaning bricks or concrete.
The difference between the two is that muriatic acid is a strong acid and vinegar is a weak acid. Muriatic acid is strong because it is very good at transferring an H+ ion to a water molecule. In a 6 M solution of hydrochloric acid, 99.996% of the HCl molecules react with water to form H3O+ and Cl- ions.

Vinegar is a weak acid because it is not very good at transferring H+ ions to water. In a 1 M solution, less than 0.4% of the CH3CO2H molecules react with water to form H3O+ and CH3CO2- ions.

CH3CO2H(aq) + H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gif H3O+(aq) + CH3CO2-(aq)

More than 99.6% of the acetic acid molecules remain intact.



The relative strengths of acids is often described in terms of an acid-dissociation equilibrium constant, Ka. To understand the nature of this equilibrium constant, let's assume that the reaction between an acid and water can be represented by the following generic equation.

In other words,some of the HA molecules react to form H3O+ and A- ions,


.
By convention, the concentrations of these ions in units of moles per liter are represented by the symbols [H3O+] and [A-]. The concentration of the HA molecules that remain in solution is represented by the symbol [HA].
The value of Ka for acid is calculated from the following equation.

http://chemed.chem.purdue.edu/genche...s/11_s9eq1.gif

When a strong acid dissolves in water, the acid reacts extensively with water to form H3O+ and A- ions. (Only a small residual concentration of the HA molecules remains in solution.) The product of the concentrations of the H3O+ and A- ions is therefore much larger than the concentration of the HA molecules, so Ka for a strong acid is greater than 1.
Example: Hydrochloric acid has a Ka of roughly 1 x 106.

http://chemed.chem.purdue.edu/genche...s/11_s9eq2.gif

Weak acids, on the other hand, react only slightly with water. The product of the concentrations of the H3O+ and A- ions is therefore smaller than the concentration of the residual HA molecules. As a result, Ka for a weak acid is less than 1.
Example: Acetic acid has a Ka of only 1.8 x 10-5.

http://chemed.chem.purdue.edu/genche...s/11_s9eq3.gif

Ka can therefore be used to distinguish between strong acids and weak acids.

Strong acids: Ka > 1Weak acids: Ka < 1

The Relative Strengths of Conjugate Acid-base Pairs

• Strong acids have a weak conjugate base.

Example: HCl is a strong acid. If HCl is a strong acid, it must be a good proton donor. HCl can only be a good proton donor, however, if the Cl- ion is a poor proton acceptor. Thus, the Cl- ion must be a weak base.

HCl(g) +H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gifH3O+(aq) + Cl-(aq) Strong
acid Weak
base

• Strong bases have a weak conjugate acid.

Example: Let's consider the relationship between the strength of the ammonium (NH4+) and its conjugate base, ammonia (NH3). The NH4+ ion is a weak acid because ammonia is a reasonably good base.

NH4+(aq)+H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gifH3O+(aq)+NH3(aq)Weak
acid Good
base



Comparing Relative Strengths of Pairs of Acids and Bases

The value of Ka for an acid can be used to decide whether it is a strong acid or a weak acid, in an absolute sense. It can also be used l to compare the relative strengths of a pair of acids.
Example: Consider HCl and the H3O+ ion.

HCl Ka = 1 x 106H3O+ Ka = 55

These Ka values suggest that both are strong acids, but HCl is a stronger acid than the H3O+ ion.
A high proportion of the HCl molecules in an aqueous solution reacts with water to form H3O+ and Cl- ions. The Brnsted theory suggests that every acid-base reaction converts an acid into its conjugate base and a base into its conjugate acid.

http://chemed.chem.purdue.edu/genche...cs/11_s11a.gif

There are two acids and two bases in this reaction. The stronger acid, however, is on the left side of the equation.

HCl(g)+H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gifH3O+(aq)+ Cl-(aq) stronger
acid weaker
acid

The general rules suggest that the stronger of a pair of acids must form the weaker of a pair of conjugate bases. The fact that HCl is a stronger acid than the H3O+ ion implies that the Cl- ion is a weaker base than water.

Acid strength: HCl > H3O+Base strength: Cl- < H2O

Thus, the equation for the reaction between HCl and water can be written as follows.

HCl(g)+H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gifH3O+(aq)+Cl-(aq) stronger
acid stronger
base weaker
acid weaker
base

It isn't surprising that 99.996% of the HCl molecules in a 6 M solution react with water to give H3O+ ions and Cl- ions. The stronger of a pair of acids should react with the stronger of a pair of bases to form a weaker acid and a weaker base.
Let's look at the relative strengths of acetic acid and the H3O+ ion.

CH3CO2H Ka = 1.8 x 10-5H3O+ Ka = 55

The values of Ka for these acids suggest that acetic acid is a much weaker acid than the H3O+ ion, which explains why acetic acid is a weak acid in water. Once again, the reaction between the acid and water must convert the acid into its conjugate base and the base into its conjugate acid.

http://chemed.chem.purdue.edu/genche...cs/11_s11b.gif

But this time, the stronger acid and the stronger base are on the right side of the equation.

CH3CO2H(aq)+H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gifH3O+(aq) +CH3CO2-(aq)weaker
acid weaker
base stronger
acid stronger
base

As a result, only a few of the CH3CO2H molecules actually donate an H+ ion to a water molecule to form the H3O+ and CH3CO2- ions.


The magnitude of Ka can also be used to explain why some compounds that qualify as Brnsted acids or bases don't act like acids or bases when they dissolve in water. When the value of Ka for an acid is relatively large, the acid reacts with water until essentially all of the acid molecules have been consumed. Sulfuric acid (Ka = 1 x 103), for example, reacts with water until 99.9% of the H2SO4 molecules in a 1 M solution have lost a proton to form HSO4- ions.

H2SO4(aq) + H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gif H3O+(aq) + HSO4-(aq)

As Ka becomes smaller, the extent to which the acid reacts with water decreases.
As long as Ka for the acid is significantly larger than the value of Ka for water, the acid will ionize to some extent. Acetic acid, for example, reacts to some extent with water to form H3O+ and CH3CO2-, or acetate, ions.

CH3CO2H(aq) + H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gif H3O+(aq) + CH3CO2-(aq)

As the Ka value for the acid approaches the Ka for water, the compound becomes more like water in its acidity. Although it is still a Brnsted acid, it is so weak that we may be unable to detect this acidity in aqueous solution.
Some potential Brnsted acids are so weak that their Ka values are smaller than water's. Ammonia, for example, has a Ka of only 1 x 10-33. Although NH3 can be a Brnsted acid, because it has the potential to act as a hydrogen-ion donor, there is no evidence of this acidity when it dissolves in water.

The Leveling Effect of Water

All strong acids and bases seem to have the same strength when dissolved in water, regardless of the value of Ka. This phenomenon is known as the leveling effect of water http://chemed.chem.purdue.edu/genche...raphics/em.gifthe tendency of water to limit the strength of strong acids and bases. We can explain this by noting that strong acids react extensively with water to form the H3O+ ion. More than 99% of the HCl molecules in hydrochloric acid react with water to form H3O+ and Cl- ions, for example,

HCl(g) + H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gif H3O+(aq) + Cl-(aq)

and more than 99% of the H2SO4 molecules in a 1 M solution react with water to form H3O+ ions and HSO4- ions.

H2SO4(aq) + H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gif H3O+(aq) + HSO4-(aq)

Thus, the strength of strong acids is limited by the strength of the acid (H3O+) formed when water molecules pick up an H+ ion.
A similar phenomenon occurs in solutions of strong bases. Strong bases react quantitatively with water to form the OH- ion. Once this happens, the solution cannot become any more basic. The strength of strong bases is limited by the strength of the base (OH-) formed when water molecules lose an H+ ion.

The Advantages of the Brnsted Definition

The Brnsted definition of acids and bases offers many advantages over the Arrhenius and operational definitions.

• It expands the list of potential acids to include positive and negative ions, as well as neutral molecules.
• It expands the list of bases to include any molecule or ion with at least one pair of nonbonding valence electrons.
• It explains the role of water in acid-base reactions: Water accepts H+ ions from acids to form the H3O+ ion.
• It can be expanded to include solvents other than water and reactions that occur in the gas or solid phases.
• It links acids and bases into conjugate acid-base pairs.
• It can explain the relationship between the strengths of an acid and its conjugate base.
• It can explain differences in the relative strengths of a pair of acids or a pair of bases.
• It can explain the leveling effect of water http://chemed.chem.purdue.edu/genche...raphics/em.gif the fact that strong acids and bases all have the same strength when dissolved in water.

Because of these advantages, whenever chemists use the words acid or base without any further description, they are referring to a Brnsted acid or a Brnsted base.

pH As A Measure of the Concentration of the H3O+ Ion

Pure water is both a weak acid and a weak base. By itself, water forms only a very small number of the H3O+ and OH- ions that characterize aqueous solutions of stronger acids and bases.

H2O(l)+H2O(l)http://chemed.chem.purdue.edu/genche...s/equilibr.gif H3O+(aq)+OH-(aq) base acid acid base

The concentrations of the H3O+ and OH- ions in water can be determined by carefully measuring the ability of water to conduct an electric current. At 25oC, the concentrations of these ions in pure water is 1.0 x 10-7 moles per liter.

[H3O+] = [OH-] = 1.0 x 10-7 M (at 25C)

When we add a strong acid to water, the concentration of the H3O+ ion increases.

HCl(aq) + H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gif H3O+(aq) + Cl-(aq)

At the same time, the OH- ion concentration decreases because the H3O+ ions produced in this reaction neutralize some of the OH- ions in water.

H3O+(aq) + OH-(aq) http://chemed.chem.purdue.edu/genche...s/equilibr.gif 2 H2O(l)

The product of the concentrations of the H3O+ and OH- ions is constant, no matter how much acid or base is added to water. In pure water at 25oC, the product of the concentration of these ions is 1.0 x 10-14.

[H3O+][OH-] = 1.0 x 10-14

The range of concentrations of the H3O+ and OH- ions in aqueous solution is so large that it is difficult to work with. In 1909 the Danish biochemist S. P. L. Sorenson suggested reporting the concentration of the H3O+ ion on a logarithmic scale, which he named the pH scale. Because the H3O+ ion concentration in water is almost always smaller than 1, the log of these concentrations is a negative number. To avoid having to constantly work with negative numbers, Sorenson defined pH as the negative of the log of the H3O+ ion concentration.

pH = -log [H3O+]

The concept of pH compresses the range of H3O+ ion concentrations into a scale that is much easier to handle. As the H3O+ ion concentration decreases from roughly 100 to 10-14, the pH of the solution increases from 0 to 14.
If the concentration of the H3O+ ion in pure water at 25oC is 1.0 x 10-7 M, the pH of pure water is 7.

pH = -log [H3O+] = -log (1.0 x 10-7) = 7

When the pH of a solution is less than 7, the solution is acidic. When the pH is more than 7, the solution is basic.

Acidic: pH < 7 Basic: pH > 7

pH of Common Acids and Bases

The pH of a solution depends on the strength of the acid or base in the solution. Measurements of the pH of dilute solutions are therefore good indicators of the relative strengths of acids and bases. Values of the pH of 0.10 M solutions of a number of common acids and bases are given in the table below.

pH of 0.10 M Solutions of Common Acids and Bases
Compound pHHCl (hydrochloric acid) 1.1H2SO4 (sulfuric acid) 1.2NaHSO4 (sodium hydrogen sulfate) 1.4H2SO3 (sulfurous acid) 1.5H3PO4 (phosphoric acid) 1.5HF (hydrofluoric acid) 2.1CH3CO2H (acetic acid) 2.9H2CO3 (carbonic acid) 3.8 (saturated solution)H2S (hydrogen sulfide) 4.1NaH2PO4 (sodium dihydrogen phosphate) 4.4NH4Cl (ammonium chloride) 4.6HCN (hydrocyanic acid) 5.1Na2SO4 (sodium sulfate) 6.1NaCl (sodium chloride) 6.4NaCH3CO2 (sodium acetate) 8.4NaHCO3 (sodium bicarbonate) 8.4Na2HPO4 (sodium hydrogen phosphate) 9.3Na2SO3 (sodium sulfite) 9.8NaCN (sodium cyanide) 11.0NH3 (aqueous ammonia) 11.1Na2CO3 (sodium carbonate) 11.6Na3PO4 (sodium phosphate) 12.0NaOH (sodium hydroxide, lye) 13.0

When all other factors are kept constant, acids become stronger as the Xhttp://chemed.chem.purdue.edu/genche...raphics/em.gifH bond becomes more polar. The second-row nonmetal hydrides, for example, become more acidic as the difference between the electronegativity of the X and H atoms increases. HF is the strongest of these four acids, and CH4 is one of the weakest Brnsted acids known.

When these compounds act as an acid, an H-X bond is broken to form H+ and X- ions. The more polar this bond, the easier it is to form these ions. Thus, the more polar the bond, the stronger the acid.
An 0.1 M HF solution is moderately acidic. Water is much less acidic, and the acidity of ammonia is so small that the chemistry of aqueous solutions of this compound is dominated by its ability to act as a base.

At first glance, we might expect that HF, HCl, HBr, and HI would become weaker acids as we go down this column of the periodic table because the X-H bond becomes less polar. Experimentally, we find the opposite trend. These acids actually become stronger as we go down this column.
This occurs because the size of the X atom influences the acidity of the X-H bond. Acids become stronger as the X-H bond becomes weaker, and bonds generally become weaker as the atoms get larger as shown in the figure below.

The Ka data for HF, HCl, HBr, and HI reflect the fact that the X-H bond-dissociation enthalpy (BDE) becomes smaller as the X atom becomes larger.

The charge on a molecule or ion can influence its ability to act as an acid or a base. This is clearly shown when the pH of 0.1 M solutions of H3PO4 and the H2PO4-, HPO42-, and PO43- ions are compared.

Compounds become less acidic and more basic as the negative charge increases.

There is no difference in the polarity, size, or charge when we compare oxyacids of the same element, such as H2SO4 and H2SO3 or HNO3 and HNO2, yet there is a significant difference in the strengths of these acids. Consider the following Ka data, for example.

The acidity of these oxyacids increases significantly as the oxidation state of the central atom becomes larger. H2SO4 is a much stronger acid than H2SO3, and HNO3 is a much stronger acid than HNO2. This trend is easiest to see in the four oxyacids of chlorine.

This factor of 1011 difference in the value of Ka for hypochlorous acid (HOCl) and perchloric acid (HOClO3) can be traced to the fact that there is only one value for the electronegativity of an element, but the tendency of an atom to draw electrons toward itself increases as the oxidation number of the atom increases.
As the oxidation number of the chlorine atom increases, the atom becomes more electronegative. This tends to draw electrons away from the oxygen atoms that surround the chlorine, thereby making the oxygen atoms more electronegative as well, as shown in the figure below. As a result, the O-H bond becomes more polar, and the compound becomes more acidic.



The relative strengths of Brnsted bases can be predicted from the relative strengths of their conjugate acids combined with the general rule that the stronger of a pair of acids always has the weaker conjugate base.

For more than 300 years, substances that behaved like vinegar have been classified as acids, while those that have properties like the ash from a wood fire have been called alkalies or bases. The name "acid" comes from the Latin acidus, which means "sour," and refers to the sharp odor and sour taste of many acids. Vinegar tastes sour because it is a dilute solution of acetic acid in water; lemon juice is sour because it contains citric acid; milk turns sour when it spoils because of the formation of lactic acid; and the sour odor of rotten meat can be attributed to carboxylic acids such as butyric acid formed when fat spoils.
Today, when chemists use the words "acid" or "base" they refer to a model developed independently by Br&oslash;nsted, Lowry, and Bjerrum. Since the most explicit statement of this theory was contained in the writings of Br&oslash;nsted, it is most commonly known as the "Br&oslash;nsted acid-base" theory.

Br&oslash;nsted argued that all acid-base reactions involve the transfer of an H+ ion, or proton. Water reacts with itself, for example, by transferring an H+ ion from one molecule to another to form an H3O+ ion and an OH- ion.

According to this theory, an acid is a "proton donor" and a base is a "proton acceptor."
Acids are often divided into categories such as "strong" and "weak." One measure of the strength of an acid is the acid-dissociation equilibrium constant, Ka, for that acid.

When Ka is relatively large, we have a strong acid.

When it is small, we have a weak acid.

When it is very small, we have a very weak acid.

In 1909, S. P. L. S&oslash;renson suggested that the enormous range of concentrations of the H3O+ and OH- ions in aqueous solutions could be compressed into a more manageable set of data by taking advantage of logarithmic mathematics and calculating the pH or pOH of the solution.

The "p" in pH and pOH is an operator that indicates that the negative of the logarithm should be calculated for any quantity to which it is attached. Thus, pKa is the negative of the logarithm of the acid-dissociation equilibrium constant.

The only disadvantage of using pKa as a measure of the relative strengths of acids is the fact that large numbers now describe weak acids, and small (negative) numbers describe strong acids.



An important features of the Br&oslash;nsted theory is the relationship it creates between acids and bases. Every Br&oslash;nsted acid has a conjugate base, and vice versa.

Just as the magnitude of Ka is a measure of the strength of an acid, the value of Kb reflects the strength of its conjugate base. Consider what happens when we multiply the Ka expression for a generic acid (HA) by the Kb expression for its conjugate base (A-).

If we now replace each term in this equation by the appropriate equilibrium constant, we get the following equation.

Because the product of Ka times Kb is a relatively small number, either the acid or its conjugate base can be "strong." But if one is strong, the other must be weak. Thus, a strong acid must have a weak conjugate base.

A strong base, on the other hand, must have a weak conjugate acid.

Water has a limiting effect on the strength of acids and bases. All strong acids behave the same in water -- 1 M solutions of the strong acids all behave as 1 M solutions of the H3O+ ion -- and very weak acids cannot act as acids in water. Acid-base reactions don't have to occur in water, however. When other solvents are used, the full range of acid-base strength shown in the following table can be observed.




The strongest acids are in the upper-left corner of this table; the strongest bases in the bottom-right corner. Each base is strong enough to deprotonate the acid in any line above it. The hydride ion (H-), for example, can convert an alcohol into its conjugate base

and the amide (NH2-) ion can deprotonate an alkyne.

It is easy to understand why aqueous solutions of HCl or CH3CO2H are acidic. The following data for the pH of 0.1 M solutions of transition-metal ions are a bit harder to explain.

We can't attribute the acidity of these solutions to the Cl- or NO3- ions because these ions are weak bases. The acidity of these solutions must result from the behavior of the Fe3+, Al3+, and Cu2+ ions.
The Fe3+, Al3+, and Cu2+ ions can't be Brnsted acids by themselves. They can only act as proton donors by influencing the ability of the neighboring water molecules to give up H+ ions. They do this by first forming covalent bonds to six water molecules to form a complex ion.
Water molecules covalently bound to one of these metal ions are more acidic than normal. Thus, reactions such as the following occur.

Al(H2O)63+(aq) + H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gif Al(H2O)5(OH)2+(aq) + H3O+(aq)

Fe(H2O)63+(aq) + H2O(l) http://chemed.chem.purdue.edu/genche...s/equilibr.gif Fe(H2O)5(OH)2+(aq) + H3O+(aq)
These reactions give rise to a net increase in the H3O+ ion concentration in these solutions, thereby making the solutions acidic.

In 1923 G. N. Lewis suggested another way of looking at the reaction between H+ and OH- ions. In the Brnsted model, the OH- ion is the active species in this reaction http://chemed.chem.purdue.edu/genche...raphics/em.gif it accepts an H+ ion to form a covalent bond. In the Lewis model, the H+ ion is the active specieshttp://chemed.chem.purdue.edu/genche...raphics/em.gifit accepts a pair of electrons from the OH- ion to form a covalent bond.

In the Lewis theory of acid-base reactions, bases donate pairs of electrons and acids accept pairs of electrons. A Lewis acid is therefore any substance, such as the H+ ion, that can accept a pair of nonbonding electrons. In other words, a Lewis acid is an electron-pair acceptor. A Lewis base is any substance, such as the OH- ion, that can donate a pair of nonbonding electrons. A Lewis base is therefore an electron-pair donor.
One advantage of the Lewis theory is the way it complements the model of oxidation-reduction reactions. Oxidation-reduction reactions involve a transfer of electrons from one atom to another, with a net change in the oxidation number of one or more atoms.

The Lewis theory suggests that acids react with bases to share a pair of electrons, with no change in the oxidation numbers of any atoms. Many chemical reactions can be sorted into one or the other of these classes. Either electrons are transferred from one atom to another, or the atoms come together to share a pair of electrons.
The principal advantage of the Lewis theory is the way it expands the number of acids and therefore the number of acid-base reactions. In the Lewis theory, an acid is any ion or molecule that can accept a pair of nonbonding valence electrons. In the preceding section, we concluded that Al3+ ions form bonds to six water molecules to give a complex ion.

This is an example of a Lewis acid-base reaction. The Lewis structure of water suggests that this molecule has nonbonding pairs of valence electrons and can therefore act as a Lewis base. The electron configuration of the Al3+ ion suggests that this ion has empty 3s, 3p, and 3d orbitals that can be used to hold pairs of nonbonding electrons donated by neighboring water molecules.

Thus, the Al(H2O)63+ ion is formed when an Al3+ ion acting as a Lewis acid picks up six pairs of electrons from neighboring water molecules acting as Lewis bases to give an acid-base complex, or complex ion.
The Lewis acid-base theroy explains why BF3 reacts with ammonia. BF3 is a trigonal-planar molecule because electrons can be found in only three places in the valence shell of the boron atom. As a result, the boron atom is sp2 hybridized, which leaves an empty 2pz orbital on the boron atom. BF3 can therefore act as an electron-pair acceptor, or Lewis acid. It can use the empty 2pz orbital to pick up a pair of nonbonding electrons from a Lewis base to form a covalent bond. BF3 therefore reacts with Lewis bases such as NH3 to form acid-base complexes in which all of the atoms have a filled shell of valence electrons, as shown in the figure below.
The Lewis acid-base theory can also be used to explain why nonmetal oxides such as CO2 dissolve in water to form acids, such as carbonic acid H2CO3.

In the course of this reaction, the water molecule acts as an electron-pair donor, or Lewis base. The electron-pair acceptor is the carbon atom in CO2. When the carbon atom picks up a pair of electrons from the water molecule, it no longer needs to form double bonds with both of the other oxygen atoms as shown in the figure below

One of the oxygen atoms in the intermediate formed when water is added to CO2 carries a positive charge; another carries a negative charge. After an H+ ion has been transferred from one of these oxygen atoms to the other, all of the oxygen atoms in the compound are electrically neutral. The net result of the reaction between CO2 and water is therefore carbonic acid, H2CO3.

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