Question

Classify each of these species as a Bronsted acid or base, or both: (a) \(\mathrm{H}_{2} \mathrm{O},\) (b) \(\mathrm{OH}^{-},\) (c) \(\mathrm{H}_{3} \mathrm{O}^{+}\), (d) \(\mathrm{NH}_{3}\) (e) \(\mathrm{NH}_{4}^{+},\) (f) \(\mathrm{NH}_{2}^{-},(\mathrm{g}) \mathrm{NO}_{3}^{-},\) (h) \(\mathrm{CO}_{3}^{2-}\) (i) HBr, (j) HCN.

Short answer

H2O is both a Bronsted acid and base. OH- and NH3 are Bronsted acids. H3O+, NH4+, HBr and HCN are Bronsted bases. NO3- is neither a Bronsted acid nor base.

Step-by-step solution

Classify H2O as acid or base or both

H2O can donate a H+ ion to become OH- or accept a H+ ion to become H3O+. Therefore, it acts as both a Bronsted acid and base.

Classify OH- as acid or base or both

OH- can only accept a H+ ion to become H2O. Therefore, it acts as a Bronsted acid.

Classify H3O+ as acid or base or both

H3O+ can only donate a H+ ion to become H2O. Therefore, it acts as a Bronsted base.

Classify NH3 as acid or base or both

NH3 can only accept a H+ ion to become NH4+. Therefore, it acts as a Bronsted acid.

Classify NH4+ as acid or base or both

NH4+ can only donate a H+ to become NH3. Therefore, it acts as a Bronsted base.

Classify NH2- as acid or base or both

NH2- can only accept a H+ ion to become NH3. Therefore, it acts as a Bronsted acid.

Classify NO3- as acid or base or both

NO3- cannot donate or accept a H+ ion. Therefore, it is neither a Bronsted acid nor base.

Classify CO32- as acid or base or both

CO32- can only accept a H+ ion to become HCO3-. Therefore, it acts as a Bronsted acid.

Classify HBr as acid or base or both

HBr can only donate a H+ ion to become Br-. Therefore, it acts as a Bronsted base.

Classify HCN as acid or base or both

HCN can only donate a H+ ion to become CN-. Therefore, it acts as a Bronsted base.

Key concepts

Acid-Base Classification

In the realm of chemistry, the classification of substances as acids or bases is foundational to understanding numerous reactions and processes. According to the Bronsted acid-base theory, substances can be classified depending on whether they donate or accept protons (H⁺ ions). This classification helps in predicting the behavior of acids and bases in solution.

- **Bronsted Acid**: A substance that can donate a proton to another substance. This means it releases a H⁺ ion in a chemical reaction.
- **Bronsted Base**: A substance that can accept a proton, meaning it can attract and hold a H⁺ ion during a reaction.

It's important to understand that the same compound can sometimes act as both an acid and a base, depending on the circumstances. Such substances are called amphoteric, like water ( H₂O), which can function as either a Bronsted acid or base depending on the surrounding environment.

Bronsted Acid

The concept of a Bronsted acid revolves around its ability to donate protons. This donation process is integral as it influences the acidity of a solution. When a Bronsted acid donates a H⁺ ion, it forms its conjugate base. For example, when hydrochloric acid (HCl) donates a hydrogen ion, it transforms into its conjugate base, Cl^-.

- In the exercise, substances like H₃O⁺ and NH₄⁺ are examples of Bronsted acids. They donate protons in reactions, thus showcasing their acidic properties.
- Interesting examples include NH₄⁺, which loses a H⁺ to become NH₃, and HCN, which donates a proton to form CN^-.

Remember, identifying proton donation helps classify Bronsted acids easily. Knowing this can assist in solving reactions where ion exchange occurs due to proton movement.

Bronsted Base

Bronsted bases are fascinating as they accept protons. In doing so, they effectively "neutralize" acidity by binding H⁺ ions. A Bronsted base that has accepted a proton forms its conjugate acid.

- For example, in the exercise, substances like OH^-, NH₃, and CO₃^2- are classic Bronsted bases because they can accept protons.
- When NH₃ accepts a proton, it becomes NH₄⁺, and similarly, OH^- becomes H₂O upon proton acceptance.

Identifying how substances react upon accepting protons can greatly aid in understanding the their behavior in chemical solutions. Recognizing the base's ability to form a conjugate acid explains the dynamic balance of reactions.

Acid-Base Reactions

Acid-base reactions are encounters that can drastically change the pH levels of a solution. These reactions involve the transfer of protons between the reactants. It's essential to identify which role each substance plays in chemical equations. Recognizing these roles allows predictions about reaction products and their characteristics.

- A typical example includes the reaction between HCl and NH₃, where NH₃ (a base) accepts a proton from HCl (an acid) to form NH₄⁺ and Cl^-.
- The product of such reactions often involves the formation of a conjugate base and a conjugate acid. For instance, in the exercise, when H₃O⁺ loses a proton, it forms H₂O, its conjugate base.

Understanding these reactions is key to solving complex chemical problems and predicting the behavior of mixed solutions. They illustrate the application of Bronsted theory to real-world chemical interactions.