14.1 Bronsted Lowry Acids & Bases

1. Bronsted-Lowry Acids and Bases

Image: courtesy of Robert Swarbrick Wikimedia https://en.wikipedia.org/wiki/

Lemon_battery#/media/

File:Lemon_Battery_With_LED_V2.svg

Modified by Eileen H. Kramer

A lemon can be a power source. If we place copper and zinc electrode’s in the lemon’s pulp, something in that lemon reacts with the metal to produce electricity that powers a small device.

We might guess that acid in the lemon acts as an electrolyte. We know lemons are acidic because their juice is sour, just as we know that a paste of the base, baking soda and water, is slimy. We remember from school that acids turn litmus paper red, while bases turn litmus paper blue.

In the 19th Century, Arrhenius proposed a definition of acids and bases. An acid is a molecule that dissolves in water and donates a hydrogen H⁺ or hydronium H3O⁺ ion. An Arrhenius base is a molecule that dissolves in water and which produces at lease one hydroxide OH^– ion.

The Arrhenius model has no place for acids or bases that do not dissolve well or at all in aqueous solution, yet such acids exist. Benzoic acid dissolves poorly in water due to its nonpolar, benzene ring. 

— Image courtesy of Ben Mills and Jynto https://en.wikipedia.org/wiki/Benzoic_acid#/media/File:Benzoic-acid-3D-balls.png

Modified by Eileen H. Kramer

In 1923 chemists, Johannes Nicholaus Brønsted and Thomas Martin Lowry, developed an acid/base definition based on proton donation and acceptance, and one we can use with partially polar molecules.  An acid is a chemical compound which dissociates and donates its hydrogen H⁺ or hydronium ions H3O⁺ when dissolved in water, while a base dissociates and accepts either a hydrogen H⁺ or hydronium ion H3O⁺ when in aqueous solution. Note: a Bronstead-Lowry acid is also an electron acceptor according to the Lewis model, while a Bronstead-Lowry base donates one or more electrons.

These videos below examine both the Arrhenius model, the Bronsted – Lowry Acid/Base concept, and acid – base conjugate pairs.

Let’s look at some acids and bases:

HCl is a well-known acid. In water HCl becomes H₃O⁺ and Cl^– ions.

HCl(aq)   +  H₂O(l)       H₃O ⁺(aq)        +    Cl ^–(aq)

Ammonia, NH₃, a well-known base becomes NH₄⁺ and OH^– in aqueous solution.

NH₃(aq)      +   H₂O +(aq)        NH₄ ⁺(aq)   +     OH ^– (aq)

HCl is a Bronsted – Lowry Acid, and NH3 is a Bronsted – Lowry Base.

The Bronsted – Lowry acids include their conjugate opposites. When HCl dissolves in water, the Cl^– ion is its conjugate base. When we dissolve ammonia in water, NH₄⁺ is its conjugate acid.

A conjugate acid gives up a proton (and accepts an electron(s)). The Cl^– loses a proton when the H in HCl forms H₃0⁺.  A conjugate base accepts a proton as when NH₃ becomes NH₄⁺. A conjugate base is also a Lewis electron donor.

Remember, a conjugate acid forms when an acid loses H⁺, while a conjugate base forms when a base captures H⁺.  And yes, both NH₃ and HCl both have respective conjugate acids and bases when in solution together.

HCl(aq)   +   NH₃(aq)         NH₄ +(aq)         +          Cl – (aq)

Acid                Base                  Conjugate acid              Conjugate base

Figure 14.2 illustrates this concept.

Reference: https://courses.lumenlearning.com/chemistryformajors/chapter/bronsted-lowry-acids-and-bases-2/

The Bronsted – Lowry acid/base concept is considered as an improvement to the Arrhenius acid – base concept. The Arrhenius acid – base concept in 1884 defined an acid as a chemical compound that is dissolved in water and produces protons H ⁺ cations or hydronium ions. A base on the other hand was defined as a chemical compound that is dissolved in water and produces the hydroxide OH ^– anions.

The videos below illustrate the concept of Arrhenius and Bronsted – Lowry Acid/Base concepts as well as the acid – base conjugate pairs concepts:

When acids and bases (and also ionic salts) dissociate in water they ionize.  Water can react as a base or as an acid depending on compounds dissolved in it. This characteristic of water is called “Amphotery” and water is amphoteric which means it can act as either an acid or base.

For example:

NH₃(aq)       +      H₂O(l)       NH4 ⁺(aq)       +        OH ^– (aq)

Base                     Acid                  Conjugate Acid           Conjugate Base

HCl(aq)      +      H₂O(l)       H₃O ⁺(aq)       +        Cl ^–(aq)

Acid                    Base                     Conjugate Acid          Conjugate Base

Water is not the only amphoteric compound. HCO₃^– , bicarbonate or hydrogen bicarbonate ion can work as either a base or acid.

H(CO₃)^– (aq) +   H2O           H₃O⁺(aq)       +        CO₃^-2(aq)

Acid                  Base                    Conjugate Acid        Conjugate Base

H(CO₃)^– (aq)    +   HCl (aq)      H₂(CO₃) (aq)       +    Cl^– (aq)

Base                           Acid                 Conjugate Acid      Conjugate Base

When the same molecules react together, with one being an acid and the other being a base, the molecules autoionize.

H₂O(l)      +     H₂O(l)          H₃O⁺(aq)       +      OH^– (aq)

Acid                 Base                 Conjugate Acid       Conjugate Base

Waters dissociation constant,  K_w quantitatively describes H₂O’s autoinonzation.

K_w   =  [ H₃O ⁺(aq) ] x [ OH ^– (aq) ]  /  [ H₂O(l) ] x [ H₂O(l) ]

or

K_w = [ H₃O⁺(aq) ] x [ OH ^– (aq) ]  /[H₂O(l)]² 

This is specific version of the acid dissociation constant K_a.

Ka = [product1][product2]/[reactant]²

Because liquid water in the K_w equation does not change concentration, we can discard K_w‘s denominator and simplify to

K_w = [H₃O⁺]*[OH^–]

Scientists learned experimentally that

[ H₃O⁺(aq) ] = [ OH^– (aq) ] = 1.0 * 10^-7M at 25^o C.

Taken together [H₃O⁺]*[OH^–] = 1.0 * 10^-14M at 25^o C

Table 12.1 shows that as the temperature decreases equilibrium for water’s autoionization against producing, though the actual differences are a very large negative powers of ten and thus very small.

Reference: https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Acids_and_Bases/Acids_and_Bases_in_Aqueous_Solutions/The_pH_Scale/Temperature_Dependence_of_the_pH_of_pure_Water

At higher temperatures, the pH and pOH of pure water will decrease.

Example:

Calculate the hydronium and the hydroxide ions concentration of water at 100 ⁰C. Kw is 51.3 x 10 ^{– 14} (from the table).

Kw = 51.3 x 10 ^{– 14} = [ H₃O ⁺(aq) ] x [ OH ^– (aq) ] =  [ H₃O ⁺(aq) ] ²

[ H₃O ⁺(aq) ]  = [ OH ^– (aq) ] = √ ( 51.3 x 10 ^{– 14} )  = 7.16 x 10  ^-7 M

The relationships of Kw, H₃O ⁺ and OH ^– are interconnected:

K_w   =  [ H₃O ⁺(aq) ] x [ OH ^– (aq) ]

[ H₃O ⁺(aq) ]   =  Kw / [ OH ^– (aq) ]

[ OH ^– (aq) ]  =  Kw / [ H₃O ⁺(aq) ]

Example:

Calculate the concentration of hydroxide ion [ OH ^– (aq) ] for pure water at 25 ^oC when the [ H₃O ⁺(aq) ] is measured to be 7.12 x 10 ^-5 M

Kw = 1.0 x 10 ^-14 at 25 ^oC

[ OH ^– (aq) ]  =  Kw / [ H₃O ⁺(aq) ] =  [ 1.0 x 10 ^{– 14} ] / [ 7.12 x 10 ^-5 ] = 1.40 x 10 ^{– 10}

The videos below explain the concept of Kw, the amphoteric character of water as well as [ OH ^– (aq) ] and [ H₃O ⁺(aq) ]

https://www.youtube.com/watch?v=_vAMJzCYMN0