Chemistry Notes - Acid-Base Reactions and Properties

The instructor graded quizzes over the weekend and found some inconsistencies in the questions.
To address these inconsistencies, a half quiz will be given at the end of the session to ensure fair assessment.
Students are encouraged to ask questions about quiz material or any other topics covered over the weekend.

The topic of discussion revolves around acid-base definitions.
Arrhenius Definition:
- Acids: Substances that produce H$^+$ ions (protons) in solution.
- Bases: Substances that produce OH$^-$ ions in solution.
Bronsted-Lowry Definition:
- Acid: A proton donor.
- Base: A proton acceptor (H$^+$).

Ammonia and Water Reaction:
- Reaction: NH₃ (ammonia) + H₂O → NH₄$^+$ + OH$^-$
- Water donates a proton to ammonia, acting as an acid, and ammonia acts as a base.
Hydrochloric Acid and Water Reaction:
- Reaction: HCl + H₂O → H₃O$^+$ + Cl$^-$
- In this case, HCl acts as an acid, and water acts as a base. The product is hydronium ion (H₃O$^+$).
Water Acting as Acid and Base:
- Water can donate a proton to another water molecule: H₂O + H₂O ⇌ H₃O$^+$ + OH$^-$
- One water molecule acts as an acid while the other acts as a base.
- This highlights the amphoteric nature of water, meaning it can act as both an acid and a base.

Definition of Amphoteric:
- A substance that can act as both an acid and a base under different conditions.
Example:
- Water (H₂O) is the most common amphoteric substance.
- Aluminum (Al) is another example, reacting with both acids and bases depending on the context.

The reaction of water forming H$^+$ and OH$^-$ ions underlies the pH scale.
Neutral pH (7):
- Indicates equal concentrations of H$^+$ and OH$^-$ ions: 1 x 10$^{-7}$ moles per liter each.

Strong Acids:
- Dissociate completely in water into H$^+$ and their corresponding anions.
- Examples include:
- Hydrochloric acid (HCl)
- Hydrobromic acid (HBr)
- Hydroiodic acid (HI)
- Nitric acid (HNO₃)
- Sulfuric acid (H₂SO₄)
- Perchloric acid (HClO₄)
Weak Acids:
- Do not fully dissociate in solution and exist in equilibrium between the undissociated acid and its ions.
- Example: Hydrofluoric acid (HF)
- Despite being labeled a weak acid, HF is highly reactive and can dissolve glass.

Strong Bases:
- Generally soluble hydroxides from Group 1 (Li, Na, K, Rb, Cs) and Group 2 (Ca, Sr, Ba).
Weak Bases & Electrolytes:
- Weak bases are typically nitrogen-containing compounds and are weak electrolytes, meaning they only partially dissociate in solution.
- Example: Ammonia (NH₃) dissolves in water but only produces a small amount of NH₄$^+$ and OH$^-$.

Electrical conductivity in solutions can vary:
- Strong Acids/Bases: High conductivity due to complete ionization in solution, resulting in a bright light in conductivity tests.
- Weak Acids/Bases: Low conductivity, resulting in dim light during tests due to low ionization levels.

Certain reactions with water can produce dangerous gases or toxic substances:
- Carbonate ions (CO₃$^{2-}$) react with acids to form carbonic acid (H₂CO₃), which decomposes into water and carbon dioxide (CO₂).
- Sulfite ions (SO₃$^{2-}$) react with strong acids to produce sulfuric acid and can release sulfur dioxide (SO₂), a toxic gas.
- Hydrogen sulfide (H₂S) is a dangerous gas with a rotten egg smell and can sedate the nervous system.
- Cyanide (CN$^-$) is extremely toxic and reacts vigorously with strong acids to produce hydrogen cyanide.

A last quiz question will be given to address past confusion on material.
Students should focus on understanding and memorizing key concepts for success in assessments.