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.
<U>

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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]H3O+(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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]H3O+(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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]NH4+(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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]H3O+(aq)+ Cl-(aq)

and then losing these H+ ions to the base.

NH3(aq)+H3O+(aq)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]NH4+(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 NH3[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]NH4++ 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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]NH4Cl(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
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.


comes from the Latin stems meaning "joined together" and refers to
things that are joined, particularly in pairs, such as Brnsted acids
and bases.
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" [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] a 6 M solution of hydrochloric acid HCl(aq)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]to 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. <blockquote></blockquote>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.


<b>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.

[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]

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.

[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]

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.

[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]

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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]H3O+(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) [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]H3O+(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.

[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]

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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]H3O+(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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]H3O+(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.

[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]

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

CH3CO2H(aq)+H2O(l)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]H3O+(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) [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] 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.<blockquote>

CH3CO2H(aq) + H2O(l) [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] H3O+(aq) + CH3CO2-(aq)

</blockquote>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 [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة]the
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,<blockquote>

HCl(g) + H2O(l) [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] H3O+(aq) + Cl-(aq)

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

H2SO4(aq) + H2O(l) [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] H3O+(aq) + HSO4-(aq)

</blockquote>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 [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] 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)[ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] 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) [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] 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) [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] 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.<blockquote>

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

</blockquote>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 [center]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

[/center]
</b>

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. <blockquote></blockquote>Fe(H2O)63+(aq) + H2O(l) [ندعوك للتسجيل في المنتدى أو التعريف بنفسك لمعاينة هذه الصورة] 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.