Learning Objectives
- Identify acids, bases, and conjugate acid-base pairs according to the Brønsted-Lowry definition
- Write equations for acid and base ionization reactions
- Use the ion-product constant for water to calculate hydronium and hydroxide ion concentrations
- Describe the acid-base behavior of amphiprotic substances
The Arrhenius definition of an acid as a compound that dissolves in water to yield hydronium ions (H3O+) and a base as a compound that dissolves in water to yield hydroxide ions (\(\ce{OH-}\)). This definition is not wrong; it is simply limited. We extended the definition of an acid or a base using the more general definition proposed in 1923 by the Danish chemist Johannes Brønsted and the English chemist Thomas Lowry. Their definition centers on the proton, \(\ce{H^+}\). A proton is what remains when a normal hydrogen atom, \(\ce{^1_1H}\), loses an electron. A compound that donates a proton to another compound is called a Brønsted-Lowry acid, and a compound that accepts a proton is called a Brønsted-Lowry base. An acid-base reaction is the transfer of a proton from a proton donor (acid) to a proton acceptor (base). In a subsequent chapter of this text we will introduce the most general model of acid-base behavior introduced by the American chemist G. N. Lewis.
Brønsted-Lowry Defintions
- A compound that donates a proton to another compound is called a Brønsted-Lowry acid.
- A compound that accepts a proton is called a Brønsted-Lowry base.
Definitions of Acids and Bases: https://youtu.be/r8reN0CSIHw
Acids may be compounds such as HCl or H2SO4, organic acids like acetic acid (\(\ce{CH_3COOH}\)) or ascorbic acid (vitamin C), or H2O. Anions (such as \(\ce{HSO_4^-}\), \(\ce{H_2PO_4^-}\), \(\ce{HS^-}\), and \(\ce{HCO_3^-}\)) and cations (such as \(\ce{H_3O^+}\), \(\ce{NH_4^+}\), and \(\ce{[Al(H_2O)_6]^{3+}}\)) may also act as acids. Bases fall into the same three categories. Bases may be neutral molecules (such as \(\ce{H_2O}\), \(\ce{NH_3}\), and \(\ce{CH_3NH_2}\)), anions (such as \(\ce{OH^-}\), \(\ce{HS^-}\), \(\ce{HCO_3^-}\), \(\ce{CO_3^{2−}}\), \(\ce{F^-}\), and \(\ce{PO_4^{3−}}\)), or cations (such as \(\ce{[Al(H_2O)_5OH]^{2+}}\)). The most familiar bases are ionic compounds such as \(\ce{NaOH}\) and \(\ce{Ca(OH)_2}\), which contain the hydroxide ion, \(\ce{OH^-}\). The hydroxide ion in these compounds accepts a proton from acids to form water:
\[\ce{H^+ + OH^- \rightarrow H_2O} \label{14.2.1}\]
We call the product that remains after an acid donates a proton the conjugate base of the acid. This species is a base because it can accept a proton (to re-form the acid):
\[\text{acid} \rightleftharpoons \text{proton} + \text{conjugate base}\label{14.2.2a}\]
\[\ce{HF \rightleftharpoons H^+ + F^-} \label{14.2.2b} \nonumber\]
\[\ce{H_2SO_4 \rightleftharpoons H^+ + HSO_4^{−}}\label{14.2.2c} \nonumber\]
\[\ce{H_2O \rightleftharpoons H^+ + OH^-}\label{14.2.2d} \nonumber\]
\[\ce{HSO_4^- \rightleftharpoons H^+ + SO_4^{2−}}\label{14.2.2e} \nonumber\]
\[\ce{NH_4^+ \rightleftharpoons H^+ + NH_3} \label{14.2.2f} \nonumber\]
We call the product that results when a base accepts a proton the base’s conjugate acid. This species is an acid because it can give up a proton (and thus re-form the base):
\[\text{base} + \text{proton} \rightleftharpoons \text{conjugate acid} \label{14.2.3a}\]
\[\ce{OH^- +H^+ \rightleftharpoons H2O}\label{14.2.3b} \nonumber\]
\[\ce{H_2O + H^+ \rightleftharpoons H3O+}\label{14.2.3c} \nonumber\]
\[\ce{NH_3 +H^+ \rightleftharpoons NH4+}\label{14.2.3d} \nonumber\]
\[\ce{S^{2-} +H^+ \rightleftharpoons HS-}\label{14.2.3e} \nonumber\]
\[\ce{CO_3^{2-} +H^+ \rightleftharpoons HCO3-}\label{14.2.3f} \nonumber\]
\[\ce{F^- +H^+ \rightleftharpoons HF} \label{14.2.3g} \nonumber\]
In these two sets of equations, the behaviors of acids as proton donors and bases as proton acceptors are represented in isolation. In reality, all acid-base reactions involve the transfer of protons between acids and bases. For example, consider the acid-base reaction that takes place when ammonia is dissolved in water. A water molecule (functioning as an acid) transfers a proton to an ammonia molecule (functioning as a base), yielding the conjugate base of water, \(\ce{OH^-}\), and the conjugate acid of ammonia, \(\ce{NH4+}\):
Similarly, in the reaction of acetic acid with water, acetic acid donates a proton to water, which acts as the base. In the reverse reaction, \(H_3O^+\) is the acid that donates a proton to the acetate ion, which acts as the base. Once again, we have two conjugate acid–base pairs: the parent acid and its conjugate base (\(CH_3CO_2H/CH_3CO_2^−\)) and the parent base and its conjugate acid (\(H_3O^+/H_2O\)).
In the reaction of ammonia with water to give ammonium ions and hydroxide ions, ammonia acts as a base by accepting a proton from a water molecule, which in this case means that water is acting as an acid. In the reverse reaction, an ammonium ion acts as an acid by donating a proton to a hydroxide ion, and the hydroxide ion acts as a base. The conjugate acid–base pairs for this reaction are \(NH_4^+/NH_3\) and \(H_2O/OH^−\).
Conjugate Acid-Base Pairs: https://youtu.be/pPrp3xEQef4
Amphiprotic Species
Like water, many molecules and ions may either gain or lose a proton under the appropriate conditions. Such species are said to be amphiprotic. Another term used to describe such species is amphoteric, which is a more general term for a species that may act either as an acid or a base by any definition (not just the Brønsted-Lowry one). Consider for example the bicarbonate ion, which may either donate or accept a proton as shown here:
\[\ce{HCO^{-}3(aq) + H_2O(l) <=> CO^{2-}3(aq) + H_3O^{+}(aq)} \nonumber\]
\[ \ce{HCO^{-}3(aq) + H_2O(l) <=> H2CO3(aq) + OH^{-}(aq)} \nonumber\]
Water is the most important amphiprotic species. It can form both the hydronium ion, \(\ce{H3O^{+}}\), and the hydroxide ion, \(\ce{OH^-}\) when it undergoes autoionization:
\[\ce{2 H_2O}_{(l)} \rightleftharpoons \ce{H_3O^+}(aq)+\ce{OH^-} (aq) \nonumber\]
Example \(\PageIndex{1}\): The Acid-Base Behavior of an Amphoteric Substance
Write separate equations representing the reaction of \(\ce{HSO3-}\)
- as an acid with \(\ce{OH^-}\)
- as a base with \(\ce{HI}\)
Solution
- \(\ce{HSO3^{-}(aq) + OH^{-}(aq) <=> SO3^{2-}(aq) + H2O(l)} \)
- \(\ce{HSO^{-}3(aq) + HI(aq) <=> H2SO3(aq) + I^{-}(aq)}\)
Exercise \(\PageIndex{1}\)
Write separate equations representing the reaction of \(\ce{H2PO4-}\)
- as a base with \(\ce{HBr}\)
- as an acid with \(\ce{OH^-}\)
- Answer a
-
\[\ce{H2PO4- (aq) + HBr(aq) <=> H3PO4(aq) + Br-(aq)} \nonumber\]
- Answer b
-
\[\ce{H2PO4-}(aq)+\ce{OH^-} (aq)\rightleftharpoons \ce{HPO4^2-}(aq)+ \ce{H_2O}_{(l)} \nonumber \]
Key Concepts and Summary
A compound that can donate a proton (a hydrogen ion) to another compound is called a Brønsted-Lowry acid. The compound that accepts the proton is called a Brønsted-Lowry base. The species remaining after a Brønsted-Lowry acid has lost a proton is the conjugate base of the acid. The species formed when a Brønsted-Lowry base gains a proton is the conjugate acid of the base. Thus, an acid-base reaction occurs when a proton is transferred from an acid to a base, with formation of the conjugate base of the reactant acid and formation of the conjugate acid of the reactant base. Amphiprotic species can act as both proton donors and proton acceptors. Water is the most important amphiprotic species. It can form both the hydronium ion, H3O+, and the hydroxide ion, \(\ce{OH^-}\) when it undergoes autoionization:
\[\ce{2 H_2O}_{(l)} \rightleftharpoons \ce{H_3O^+}(aq)+\ce{OH^-} (aq) \nonumber\]
Glossary
- acid ionization
- reaction involving the transfer of a proton from an acid to water, yielding hydronium ions and the conjugate base of the acid
- amphiprotic
- species that may either gain or lose a proton in a reaction
- amphoteric
- species that can act as either an acid or a base
- autoionization
- reaction between identical species yielding ionic products; for water, this reaction involves transfer of protons to yield hydronium and hydroxide ions
- base ionization
- reaction involving the transfer of a proton from water to a base, yielding hydroxide ions and the conjugate acid of the base
- Brønsted-Lowry acid
- proton donor
- Brønsted-Lowry base
- proton acceptor
- conjugate acid
- substance formed when a base gains a proton
- conjugate base
- substance formed when an acid loses a proton
- ion-product constant for water (Kw)
- equilibrium constant for the autoionization of water
Contributors and Attributions
Paul Flowers (University of North Carolina - Pembroke),Klaus Theopold (University of Delaware) andRichard Langley (Stephen F. Austin State University) with contributing authors.Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/85abf193-2bd...a7ac8df6@9.110).
FAQs
What is the Lowry Bronsted theory of acids and bases? ›
In the Brønsted–Lowry definition of acids and bases, an acid is a proton (H⁺) donor, and a base is a proton acceptor. When a Brønsted–Lowry acid loses a proton, a conjugate base is formed. Similarly, when a Brønsted–Lowry base gains a proton, a conjugate acid is formed.
Which is the stronger Brønsted-Lowry base F − or BR −? ›Answer and Explanation: HF is a weaker acid than HBr. So, the conjugate base of HF, i.e., F− , is a stronger base than the conjugate base of HBr, i.e., Br− . HBr is strong acid while HCN is a weaker acid.
How do you identify Brønsted-Lowry acid-base pairs? ›To determine whether a substance is an acid or a base, count the hydrogens on each substance before and after the reaction. If the number of hydrogens has decreased that substance is the acid (donates hydrogen ions). If the number of hydrogens has increased that substance is the base (accepts hydrogen ions).
How many acid-base pairs are there in a Brønsted-Lowry reaction? ›A typical Brønsted-Lowry acid-base reaction contains two conjugate acid-base pairs, as shown below.
What is a Brønsted-Lowry acid and base quizlet? ›A Bronsted-Lowry Acid is a compound that donates a proton (H+ ion). A Bronsted-Lowry Base is a compound that accepts a proton (H+ ion). Strong acid. A strong acid completely dissociates into H+ ion(s) and an anion when dissolved in water.
What is the Brønsted-Lowry definition quizlet? ›According to the Brønsted-Lowry definition, an acid can donate a hydrogen ion to another substance and a base can accept a hydrogen ion. A conjugate acid-base pair consists of two substances whose formulas differ by one hydrogen ion (H⁺) or proton.
What is an example of a Bronsted base? ›Brønsted–Lowry's acids have ionizable protons that they donate to bases. Therefore, Brønsted–Lowry's acid is generally written as HA, where H+ is the donatable proton, and A- is the anion of the acid. Examples of acids are HCl, H2SO4, HNO3, and CH3COOH.
How do you know which Bronsted base is strongest? ›The central atom Cl has the least oxidation state in HClO. Hence, HClO will be the weakest acid and the corresponding conjugate base of ClO− is the strongest Bronsted base.
How do you determine the strongest Bronsted-Lowry acid and base? ›1. The strongest acids are at the bottom left, and the strongest bases are at the top right. The conjugate base of a strong acid is a very weak base, and, conversely, the conjugate acid of a strong base is a very weak acid.
What is the difference between Lewis and Brønsted-Lowry acids and bases? ›A Bronsted-Lowry acid is a compound that can donate a hydrogen ion while a Bronsted-Lowry base is a compound that can accept a hydrogen ion. A Lewis base is a compound that can donate a pair of electrons while a Lewis acid is a compound that can accept a pair of electrons.
How do you write a Brønsted-Lowry acid base equation? ›
HCl(aq) + H2O (l) → H3O+(aq) +Cl−(aq) Using the Brønsted-Lowry theory, the reaction of ammonia and hydrochloric acid in water is represented by the following equation: NH3(aq) + HCl(aq) → NH4+(aq) + Cl−(aq) Hydrochloric acid and the chlorine ion are one conjugate acid-base pair, and the ammonium ion and ammonia are the ...
What is an example of a Brønsted-Lowry base with reaction? ›The reaction of acetic acid with a non-aqueous solution, for example ammonia, results in: C H 3 C O O H + N H 3 → C H 3 C O O N H 4 . Both equations represent examples of Bronsted-Lowry reactions.
What are examples of Brønsted-Lowry conjugate acid-base pairs? ›These two species that differ by only a proton constitute a conjugate acid–base pair. For example, in the reaction of HCl with water shown below, HCl , the parent acid, donates a proton to a water molecule, the parent base, thereby forming Cl-. Thus HCl and Cl- constitute a conjugate acid–base pair.
What is an example of a Brønsted-Lowry acid? ›HCl(g) is the proton donor and therefore a Brønsted-Lowry acid, while H2O is the proton acceptor and a Brønsted-Lowry base. These two examples show that H2O can act as both a proton donor and a proton acceptor, depending on what other substance is in the chemical reaction.
What is the Brønsted-Lowry definition _____? ›According to the Brønsted-Lowry definition, an acid is a substance that can donate a proton (H+ ion) to another molecule. A base is a substance that can accept that donated H+. After the Brønsted-Lowry acid donates its proton, it becomes the conjugate base of the acid.
Which of the following is the best definition of a Brønsted-Lowry acid? ›A Bronsted-Lowry acid is defined as a substance which donates a proton in a chemical reaction.
What is the best definition of a Bronsted Lowry base? ›A Brønsted-Lowry base is any species that is capable of accepting a proton, which requires a lone pair of electrons to bond to the H+start text, H, end text, start superscript, plus, end superscript. Water is amphoteric, which means it can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base.
How does the brønsted Lowry definition of acids and bases differ from the Arrhenius definition of acids and bases? ›An Arrhenius Acid is something that donates a proton to water, and Bronsted-Lowry Concept extends this to any substance, where an acid is a proton donor and a base is a proton acceptor.
Which two substances are Brønsted bases? ›So CO32− and H2O are bronsted bases.
What are strong Bronsted acid examples? ›Strong Brønsted acids such as HCl, HBr, HI & H2SO4, rapidly add to the C=C functional group of alkenes to give products in which new covalent bonds are formed to hydrogen and to the conjugate base of the acid.
Which Brønsted-Lowry acid has the strongest base? ›
Hence, (c) phenol is the strongest bronsted acid.
Which is most Bronsted acid? ›The Structure of the Strongest Brønsted Acid: The Carborane Acid H(CHB11Cl11)
How can you tell if a Bronsted acid is strong? ›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.
What is the definition of the Brønsted-Lowry model for acids? ›A Bronsted-Lowry acid is defined as a substance that gives up or donates hydrogen ions during a chemical reaction. In contrast, aBronsted-Lowry base accepts hydrogen ions. Another way of looking at it is that a Bronsted-Lowry acid donates protons, while the base accepts protons.
What is the theory of acids and bases? ›Swedish Svante Arrhenius, in 1884 proposed the concept of acid and base based on the theory of ionization. According to Arrhenius, the acids are the hydrogen-containing compounds which give H+ ions or protons on dissociation in water and bases are the hydroxide compounds which give OH− ions on dissociation in water.
What is the importance of Bronsted acid base definition? ›An important features of the Brønsted theory is the relationship it creates between acids and bases. Every Brønsted 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.
What are acids and bases explain with examples? ›An acid is a proton donor. While a base is a proton acceptor. Acetic acid (CH3COOH) and sulphuric acid are two examples of Acid. Sodium Hydroxide (NaOH) and Ammonia are two examples of Bases.
What are the three types of acids and bases? ›There are 3 types of acids and bases: Arrhenius, Brønsted, and Lewis. Arrhenius acid dissolves in water to release H ions, and bases release OH- ions. Brønsted acids are compounds capable of donating a proton H . Brønsted bases can accept a proton.
What are the limitations of the Brønsted-Lowry theory? ›In practice, the reaction between acidic oxides and basic oxides is feasible in absence of solvent to form salt. Another limitation of Bronsted Lowery theory includes inability to explain acidic properties of some specific compounds which do not have any proton in their chemical formula but still act as an acid.
What are the advantages of Brønsted-Lowry concept of acid and base? ›Major Advantages of Brønsted-Lowry Theory
It explains the basic property of substances that do not contain hydroxide ions. It expands the role of water in acid-base reactions as more than just a solvent. It can be expanded to include solvents other than water and reactions that occur in non-aqueous states.
What is a strong Brønsted base example? ›
With the increase in the number of oxygen atoms in the given conjugate bases, the delocalization of the π bond becomes more extended. This results in decrease in electron density. Consequently, proton attraction and basicity also decreases in the order ClO−>ClO2−>ClO3−>ClO4−. Hence, the strongest Bronsted base is ClO−.
What is an example of a Brønsted-Lowry base? ›Some examples of Brønsted–Lowry bases are acetate (CH3COO–), phosphate [(PO4)3-], carbonate (CO32-), sulfide (S2-), and halide (X–). Because of its ability to both accept and donate protons, water is known as an amphoteric or amphiprotic substance.