Brpnsted-Lowry acid-base definition

Knowledge Required (1) The interpretation of acid equilibrium constants, K. (2) Meaning of terms associated with weak and strong acids and bases. (3) Brpnsted-Lowry acid-base definition. 

How are the Arrhenius and Brpnsted-Lowry acid-base definitions different How are they similar Name two Brpnsted-Lowry bases that are not Arrhenius bases. Can you do the same for acids Explain. 

The hydrogen ion accepts the lone pair of electrons from the ammonia to form the ammonium ion. The hydrogen ion, because it accepts a pair of electrons, is the Lewis acid. The ammonia, because it donates a pair of electrons, is the Lewis base. This reaction is also a Brpnsted-Lowry acid-base reaction. This illustrates that a substance may be an acid or a base by more than one definition. All Brpnsted-Lowry acids are Lewis acids, and all Brpnsted-Lowry bases are Lewis bases. However, the reverse is not necessarily true. 

You have reviewed the Brpnsted-Lowry definition of acids and bases and the meanings of pH and p a- You have learned to identify the most acidic hydrogen atoms in a molecule based on a comparison of p a values. You will see in many cases that Brpnsted-Lowry acid-base reactions either initiate or complete an organic reaction, or prepare an organic molecule for further reaction. The Lewis definition of acids and bases may have been new to you. However, you will see over and over again that Lewis acid-base reactions which involve either the donation of an electron pair to form a new covalent bond or the departure of an electron pair to break a covalent bond are central steps in many organic reactions. The vast majority of organic reactions you will study are either Brpnsted-Lowry or Lewis acid-base reactions. 

In 1923, Brpnsted and Lowry defined acids and bases on the basis of the transfer of protons. A Brpnsted-Lowry acid is any species that can donate a proton, and a Brpnsted-Lowry base is any species that can accept a proton. These definitions also include all the Arrhenius acids and bases because compounds that dissociate to give H30+ are proton donors, and compounds that dissociate to give OH are proton acceptors. (Hydroxide ion accepts a proton to form H20.)... 

The next theory of acids and bases is called the Brpnsted-Lowry Theory, proposed by Danish chemist Johannes Brpnsted and the English chemist Thomas Lowry, independently, in 1923. The definition of an acid, in this theory, sounds essentially the same as the Arrhenius acid. ABr0nsted-Lowry acid is defined as a substance that donates a proton to another species. Now, you might say, A proton is a hydrogen ion (H+), so what is the difference between the two definitions for acids Most notably, this definition doesn t require that the acid be in an aqueous solution. So a substance that donates protons, even when in a solid or vapor phase, is still acting as a Brpnsted-Lowry acid. 

Another significant difference between definitions is that Brpnsted-Lowry acids andbases need not be molecular substances. There are a variety of reactions in which ions donate or accept protons. In the sample below, note how the cyanide ion (CN ) acts as a base by accepting a proton and the bicarbonate ion (HCO,) acts as an acid by donating a proton. 

By focusing on where the proton comes from and goes to, the Brpnsted-Lowry concept expands the definition of a base to encompass a host of species that the Arrhenius definition excludes a base is any species that accepts a proton to do so, the base must have a lone electron pair. (The lone electron pair also plays the central role in the Lewis acid-base definition, as you ll see later in this chapter.)... 

The final acid-base concept we consider was developed by Gilbert N. Lewis, whose contribution to understanding the importance of valence electron pairs in molecular bonding we discussed in Chapter 9. Whereas the Brpnsted-Lowry concept focuses on the proton in defining a species as an acid or a base, the Lewis concept highlights the role of the electron pair. The Lewis acid-base definition holds that... 

A Lewis acid is an electron pair acceptor while a Brpnsted-Lowry acid is a proton donor. The proton of a Brpnsted-Lowry acid fits the definition of a Lewis acid because it accepts an electron pair when it bonds with a base. All Lewis acids are not Br0nsted-Lowry acids. A Lewis base is an electron pair donor and a Brpnsted-Lowry base is a proton acceptor. All Brpnsted-Lowry bases can be Lewis bases, and vice versa. 18.85(a) No, Zn(H20)6 ""(fl< ) + 6NH3(flf/) Zn(NH3)6"++ 6H20(/) NH3 is a weak Brpnsted-Lowry base, but a strong Lewis base. 

The Lewis definition, like the Br0nsted-Lowry definition, requires that a base have an electron pair to donate, so it does not expand the classes of bases. However, this definition greatly expands the classes of acids. Many species, such as CO2 and Cu, that do not contain H in their formula (and thus cannot be Brpnsted-Lowry acids) are Lewis acids because they accept an electron pair in reactions. Thus, the proton itself is a Lewis acid because it accepts the electron pair donated by a base ... 

The Lewis definition does not associate acidity with any particular element but rather to electronic arrangement. Lewis theory does not falsify Br0nsted-Lowry theory but extends it Brpnsted-Lowry theory is a subset of Lewis theory. All Brpnsted-Lowry bases are Lewis bases and all Brpnsted-Lowry acids are Lewis... 

In 1923, J. N. Br0nsted in Denmark and T. M. Lowry in Great Britain expanded the definition of acids and bases to include bases that do not contain OH ions. A Brpnsted-Lowry acid can donate a hydrogen ion (H ) to another substance, and a Brpnsted-Lowry base can accept a hydrogen ion (H+). 

Section 1 13 According to the Brpnsted-Lowry definitions an acid is a proton donor and a base is a proton acceptor... 

This IS a very useful relationship You should practice writing equations according to the Brpnsted-Lowry definitions of acids and bases and familiarize yourself with Table 1 7 which gives the s of various Br0n sted acids... 

At the microscopic level, the Arrhenius theory defines acids as substances which, when dissolved in water, yield the hydronium ion (H30+) or H+(aq). Bases are defined as substances which, when dissolved in water, yield the hydroxide ion (OH). Acids and bases may be strong (as in strong electrolytes), dissociating completely in water, or weak (as in weak electrolytes), partially dissociating in water. (We will see the more useful Brpnsted-Lowry definitions of acids and bases in Chapter 15.) Strong acids include ... 

Brpnsted-Lowry definition does not differ appreciably from the Arrhenius definition of hydrogen ions (acids) and hydroxide ions (bases) ... 

Section 1.17 The Lewis definitions of acids and bases provide for a more general view of acid-base reactions than either the Arrhenius or Brpnsted-Lowry picture. A Lewis acid is an electron-pair acceptor. A Lewis base is an electron-pair donor. The Lewis approach incorporates the Brpnsted-Lowry approach as a subcategory in which the atom that accepts the electron pair in the Lewis acid is a proton. 

In addition to Arrhenius acids and bases, the Brpnsted-Lowry definition includes bases that have no hydroxide ions, yet can accept protons. Consider the following examples of acids donating protons to bases. NaOH is a base under either the Arrhenius or Brpnsted-Lowry definition. The other three are Brpnsted-Lowry bases but not Arrhenius bases, because they have no hydroxide ions. 

Brpnsted-Lowry Definition Acids are proton (H1+) donors. Bases are proton (H1+) acceptors. 

Such considerations led to a more general definition of acids and bases, which was proposed independently by J. N. Brpnsted and T. M. Lowry in 1923. They defined acid as any substance (in either the molecular or the ionic state) which donates protons (H+), and a base as any substance (molecular or ionic) which accepts protons. Denoting the acid by A and the base by B, the acid-base equilibrium can be expressed as... 

In 1923, Johannes N. Brpnsted (1879-1947) and Thomas M. Lowry (1874-1936) independently defined acids and bases in a different way from the Arrhenins definitions. The resulting theory is sometimes called the Brpnsted-Lowry theory, bnt more often is referred to as just the Brpnsted theory. The Brpnsted theory extends the definitions of acid and base in a way that explains more than the Arrhenins definitions can explain. According to this theory, a Brpnsted acid is a proton donor, and a Brpnsted base is a proton acceptor. 

The Br0nsted-Lowry theory expands the definition of acids and bases to allow us to explain n ch more ol solution chemistry. For example, the Brpnsted-Lowry theory allows us to explain why a solution af ammonium nitrate tests acidic and a solution of potassium acetate tests basic. Most of the substances that we cofcider acids in the Arrhenius theory are also acids in the Brpnsted-Lowry theory, and the same is true of bases. Injboth theories, strong acids are those that react completely with water to form ions. Weak acids ionize only slightly. We can now explain this partial ionization as an equilibrium reaction of the weak acid, the ions, and the w ater. A similar statement can be made about weak bases ... 

In Chapter 2, we will concentrate on two definitions of acids and bases the Brpnsted-Lowry definition, which describes acids as proton donors and bases as proton acceptors, and the Lewis definition, which describes acids as electron pair acceptors and bases as electron pair donors. 

Lewis defined a base as an electron-pair donor and an acid as an electron-pair acceptor. This definition further expands the list to include metal ions and other electron pair acceptors as acids and provides a handy framework for nonaqueous reactions. Most of the acid-base descriptions in this book will use the Lewis definition, which encompasses the Brpnsted-Lowry and solvent system definitions. In addition to all the reactions discussed previously, the Lewis definition includes reactions such as... 

Of the numerous definitions of acids and bases that have been employed over the years, the 1923 definitions of J. N. Brpnsted and T. M. Lowry have proven to be the most useful for discussions of ionic equilibria in aqueous systems. According to the Brpnsted-Lowry model, an acid is a substance capable of donating a proton to another substance, such as water ... 

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