Question

In each of the following chemical equations, identify the conjugate acid-base pairs. a. \(\mathrm{NH}_{3}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{NH}_{4}^{+}(a q)+\mathrm{OH}^{-}(a q)\) b. \(\mathrm{PO}_{4}^{3-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{HPO}_{4}^{2-}(a q)+\mathrm{OH}^{-}(a q)\) c. \(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}(a q)+\mathrm{OH}^{-}(a q)\)

Short answer

a. Conjugate acid-base pairs: NH3/NH4+ and H2O/OH− b. Conjugate acid-base pairs: PO43-/HPO42- and H2O/OH− c. Conjugate acid-base pairs: C2H3O2−/HC2H3O2 and H2O/OH−

Step-by-step solution

a. Identify conjugate acid-base pairs#a.

In the equation NH3(aq) + H2O(l) ⇌ NH4+ (aq) + OH−(aq), consider the following comparisons: 1. NH3 and NH4+ 2. H2O and OH− NH3 and NH4+ differ by one H+ ion: NH3 + H+ → NH4+ This means that NH3 is the base and NH4+ is its conjugate acid. H2O and OH− also differ by one H+ ion: H2O - H+ → OH− This means that H2O is the acid and OH− is its conjugate base. Conjugate acid-base pairs: 1. NH3/NH4+ 2. H2O/OH−

b. Identify conjugate acid-base pairs#b.

In the equation PO43-(aq) + H2O(l) ⇌ HPO42-(aq) + OH-(aq), consider the following comparisons: 1. PO43- and HPO42- 2. H2O and OH- PO43- and HPO42- differ by one H+ ion: PO43- + H+ → HPO42- This means that PO43- is the base and HPO42- is its conjugate acid. H2O and OH− also differ by one H+ ion: H2O - H+ → OH− This means that H2O is the acid and OH− is its conjugate base. Conjugate acid-base pairs: 1. PO43-/HPO42- 2. H2O/OH−

c. Identify conjugate acid-base pairs#c.

In the equation C2H3O2−(aq) + H2O(l) ⇌ HC2H3O2(aq) + OH−(aq), consider the following comparisons: 1. C2H3O2− and HC2H3O2 2. H2O and OH− C2H3O2− and HC2H3O2 differ by one H+ ion: C2H3O2− + H+ → HC2H3O2 This means that C2H3O2− is the base and HC2H3O2 is its conjugate acid. H2O and OH− also differ by one H+ ion: H2O - H+ → OH− This means that H2O is the acid and OH− is its conjugate base. Conjugate acid-base pairs: 1. C2H3O2−/HC2H3O2 2. H2O/OH−

Key concepts

Acid-Base Reactions

Understanding acid-base reactions is essential for grasping various chemical processes. In these reactions, an acid donates a proton (H⁺) to a base. An easy way to remember is that an acid is a proton donor, while a base is a proton acceptor.

In the provided exercises, the shift of a proton from water (H₂O) to another molecule creates a conjugate acid-base pair. An important thing to note is that these reactions are reversible, which means they can move forwards and backwards, reaching a state of balance, or equilibrium. In our examples, ammonia (NH₃), phosphate (PO₄³⁻), and acetate (C₂H₃O₂⁻) act as bases and upon accepting a proton, turn into their conjugate acids: ammonium (NH₄⁺), hydrogen phosphate (HPO₄²⁻), and acetic acid (HC₂H₃O₂), respectively.

Remember:

• Acid + Base → Conjugate Acid + Conjugate Base
• The difference between members of a conjugate pair is one proton.

Grasping this concept enables you to identify and predict the behavior of substances in acid-base reactions.

Chemical Equilibrium

When we look at reversible acid-base reactions, we can't ignore the concept of chemical equilibrium. Equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction, and the concentrations of the reactants and products remain constant over time – they do not have to be equal to each other, just constant.

This state can be represented symbolically with a double-headed arrow (⇌). In the given exercise examples, reactions are shown reaching equilibrium, which means both the forward reaction (base accepting a proton becoming conjugate acid) and the reverse reaction (conjugate acid donating a proton becoming base) are occurring at the same rate. Factors like concentration, temperature, and pressure can shift this equilibrium, favoring either the products or the reactants in what is known as Le Chatelier's principle.

Understanding this will help to predict how a system will respond to changes in conditions, which is fundamental in many fields, including chemistry, biology, environmental science, and engineering.

Brønsted-Lowry Theory

The Brønsted-Lowry theory provides a more nuanced approach to acid-base reactions than the traditional Arrhenius definition. While Arrhenius focused on the formation of H⁺ and OH⁻ ions in water, Brønsted and Lowry defined acids and bases in terms of their ability to donate or accept protons (H⁺ ions) regardless of the solvent.

The exercise examples perfectly illustrate the Brønsted-Lowry theory in action. Substances like NH₃, PO₄³⁻, and C₂H₃O₂⁻ accept protons, functioning as Brønsted-Lowry bases, and in doing so, they transform into NH₄⁺, HPO₄²⁻, and HC₂H₃O₂, the corresponding conjugate acids.

Remember, the strength of an acid or base in this theory is depicted by its ability to donate or accept a proton. A strong acid will have a weak conjugate base, and a strong base will have a weak conjugate acid. The Brønsted-Lowry theory broadens our perspective on acid-base chemistry beyond aqueous solutions, making it a powerful tool for understanding a wider range of chemical reactions.