Chemistry
Brønsted-Lowry Acids and Bases
Brønsted-Lowry acids are substances that donate protons, while Brønsted-Lowry bases are substances that accept protons. This definition of acids and bases is broader than the traditional Arrhenius definition, as it includes substances that do not necessarily produce hydrogen ions in water. According to the Brønsted-Lowry theory, an acid-base reaction involves the transfer of a proton from the acid to the base.
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eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
− ) when dissolved in water. This definition only held for ionic compounds containing hydrogen or hydroxide ions and did not apply to many acids and bases that we deal with today. Since this early definition of acids and bases was so limited, two more sophisticated and general definitions of acids and bases have since been developed, which are in wide use today. These are known as the Brønsted-Lowry definition and the Lewis definition.In 1923, both J.N. Brønsted of Denmark and Thomas Lowry of England, working independently, defined an acid as a species that can donate a hydrogen ion to a base. A base was defined as a species that can accept a hydrogen ion from an acid. So, a Brønsted-Lowry acid is a proton (H+ ) donor and a Brønsted-Lowry base is a proton (H+ ) acceptor. Both acids and bases are divided into two categories, strong and weak. The Brønsted-Lowry measure of the strength of an acid is determined by the ability of the acid to give up a proton. The measure of the strength of a base is the ability of the base to attract a proton. The relative strength of an acid is often described quantitatively in terms of an acid ionization constant (K a ), which will be covered in detail in Section 5.3 .An acid-base reaction, then, involves the transfer of a proton from an acid to a base to form a new acid and a new base. When an acid loses its proton, it becomes a conjugate base and when a base gains a proton it becomes a conjugate acid . The term “conjugate” comes from Latin and means “joined together .” It particularly refers to things that are joined together in pairs. The Brønsted-Lowry acid-base reaction can be summarized by the generalized chemical equation;HA Acid+B Base→A Conjugate base+HB Conjugate acid(1)So, a conjugate acid is a species formed by the addition of a proton (H+ ) to a base. A conjugate base is a species formed by the removal of a proton from an acid. The acid and its conjugate base as well as the base and its conjugate acid are known as conjugate pairs , whose chemical formulas are related by the gain or loss of a hydrogen ion. Some examples of acids and their conjugate bases are given in Table 5.1 . Notice how the conjugate pairs in Table 5.1
eBook - ePub- Graham Patrick(Author)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
SECTION G — ACID–BASE REACTIONS
G1 Brønsted–Lowry acids and bases
Key Notes
Definition The Brønsted–Lowry definition of an acid is a molecule which can provide a proton. The Bronsted–Lowry definition of a base is a molecule which can accept that proton. Brønsted–Lowry acids A hydrogen atom attached to an electronegative atom such as a halogen, oxygen, or nitrogen is potentially acidic. Therefore, compounds containing the following functional groups (carboxylic acid, phenol, alcohol, 1° and 2° amines, and 1° and 2° amides) can act as Bronsted–Lowry acids. Brønsted–Lowry bases Examples of Bronsted–Lowry bases include negatively charged ions and neutral molecules containing oxygen or nitrogen (e.g. water, ethers, alcohols, and amines). Related topics (E3) Neutral inorganic species (E4) Organic structures (G2) Acid strength (G3) Base strength (G5) Enolates Definition
Put at its simplest, the Bronsted–Lowry definition of an acid is a molecule which can provide a proton. The Brønsted–Lowry definition of a base is a molecule which can accept that proton.An example of a simple acid/base reaction is the reaction of ammonia with water (Figure 1 ). Here, water loses a proton and is an acid. Ammonia accepts that proton and is the base.Figure 1. Reaction of ammonia with water.As far as the mechanism of the reaction is concerned, the ammonia uses its lone pair of electrons to form a new bond to the proton and is therefore acting as a nucleophile. This means that the water is acting as an electrophile.As the nitrogen uses its lone pair of electrons to form the new bond, the bond between hydrogen and oxygen must break since hydrogen is only allowed one bond. The electrons making up the O–H bond will move onto oxygen to produce a third lone pair of electrons, thus giving the oxygen a negative charge (Figure 2
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
− . The weaker acid and base of the products are favored.)The general term that covers all Brønsted–Lowry acid–base reactions is protolysis. All protolytic reactions favor the production of the (weaker, stronger) __________ acid and base.Answer: weakerArrhenius defined an acid–base neutralization as a reaction in which H2 O is always a product. Brønsted–Lowry defines an acid–base neutralization as a reaction in which the solvent is always a product, and an acid–base reaction is referred to as protolysis.The Brønsted–Lowry definition broadens our concepts of acids and bases by permitting more compounds to be considered as acids or bases that are not under the Arrhenius definition. The Brønsted–Lowry definitions are particularly useful to organic chemists, who deal quite often with nonaqueous systems. However, the Brønsted–Lowry idea still requires the presence of a proton that may be transferred from an acid to a base. This is still somewhat limiting in scope, so a third concept involving something that is present in all substances, electrons, has been developed. We now discuss that third concept.THE LEWIS ACID–BASE CONCEPT
The Arrhenius concept of acids and bases covers aqueous solutions of acids and bases. The Brønsted–Lowry concept of acids and bases covers acids and bases in other solvents as well as water. Some reactions have the characteristics of an acid–base reaction but do not undergo proton exchange.An “acid” that does not have a proton to donate would not fit the Brønsted–Lowry concept. For all practical purposes, a proton is the same as what ion? __________Answer: H+A scientist by the name of Lewis developed an acid–base theory to include those reactions that seem to behave like acid–base reactions but do not involve proton (H+
eBook - ePub- Arnost Kotyk, Jan Slavik(Authors)
- 2020(Publication Date)
- CRC Press(Publisher)
II. IONIZATION OF ACIDS AND BASES
Although water is by far the most abundant component of all living systems, its dissociation into oxonium ions is so weak that the pH of a salt solution, either extra- or intracellular, is determined by the presence of components that readily dissociate or readily bind an oxonium ion, i.e., acids and bases, respectively.Throughout modern electrochemistry, three theories of acids and bases came into prominence.The first theory, that of Arrhenius3 dates back to 1887 when he postulated the universal existence of dissociation of electrolytes in solution, supporting his views by conductometric measurements. He calculated the degree of dissociation α from the ratios of equivalent conductivities at a given and at an infinite dilution. Thus,α = Λ /It was Arrhenius who defined acids and bases in a simple way, stating that an acid (HA) is characterized by dissociation of hydrogen ionsΛ ∞
while a base (BOH) is recognized as a substance dissociating hydroxide ionsHA ⇌H ++A −(7a) BOH ⇌B ++OH −(7b) Although many of his deductions are still valid, particularly with respect to aqueous solutions, he could not foresee the behavior of acids and bases in nonaqueous solvents and the role of interactions between the solute and the solvent.The second theory is the one that is most relevant to our considerations of pH in aqueous solutions and is due to Brönsted4
eBook - ePub- Tony Farine, Mark A. Foss(Authors)
- 2013(Publication Date)
- Routledge(Publisher)
Chapter 2 , since acid–base balance is very much concerned with regulating the normal internal environment.Summary pointsAcids are substances that donate hydrogen ions during a chemical reaction.Bases are substances that accept hydrogen ions during a chemical reaction.Acids and bases are described as weak or strong, depending upon the extent of their dissociation.Acids and bases react together to produce a salt and water.The concentration of hydrogen ions is described in terms of pH.A buffer is a solution that resists a change in pH.The body’s acid–base balance is maintained by buffer systems, respiratory regulation and renal regulation.Acid–base imbalances are either acidosis or alkalosis and may be described further as respiratory or metabolic.Self-test questions (a) Citric acid. (b) Amino acid. (c) Hydrochloric acid. (d) Folic acid.1Which one of the following acids does the stomach produce?(a) Folic acid. (b) Amino acid. (c) Ascorbic acid. (d) Acetylsalicylic acid.2Which one of the following acids is not an important component of the diet?(a) An electron. (b) A proton.3Which one of the following particles may also be referred to as a hydrogen ion?
eBook - ePub- Professor Rob Beynon, J Easterby(Authors)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
Chapter 2 , but these need not concern us here.◊ Nearly all pH buffers are weak acids or bases.Notice that the weak acid can be neutral (acetic acid) or carry a positive (TrisH+ ) or negative (phosphate1– ) charge. As we develop the theory of buffers, it will become clear that these charges on the buffer species have important consequences.2. Weak acids and bases resist pH changesA buffer is able to resist changes in pH because it exists in an equilibrium between a form that has a hydrogen ion bound (conjugate acid, protonated) and a form that has lost its hydrogen ion (conjugate base, deprotonated). For the simple example of acetic acid, the equation is:CH 3COOH ⇌CH 3COO −+H +Here, the protonated form is acetic acid, with a net charge of zero, whereas the deprotonated form (acetate) has a charge of −1. The two species are in equilibrium, and this equilibrium, in common with all equilibria, can be displaced by addition of one component.Consider a solution that contains equal amounts of acetic acid and acetate ions (10 mM acetic acid, 10 mM sodium acetate, for example). If we were to add a strong acid, such as HCl, to this solution, the added H+ would displace the equilibrium to the left. Binding of H+ to CH3 COO– ‘mops up’ the added protons (Figure 3.1 ). Electrical neutrality is preserved because every H+ that reacts with a CH3 COO– anion to form the neutral CH3 COOH leaves behind a chloride (Cl– ) anion in its place. Add a strong base, such as sodium hydroxide, and the OH- ion would react with the H+ and displace the equilibrium to the right. Electrical neutrality in the solution is sustained because for every CH3 COOH that is converted to CH3 COO– , a corresponding Na+
