Classification of Chemicals: Salts, Molecules, Acids, and Bases

Salts, Molecules, Acids, & Bases

• Salts

  • Ionic compounds; contain metals and/or ammonium, often with nonmetals, or polyatomic ions.

  • Examples given: CuCl₂ (copper(II) chloride) and NH₄OH (ammonium hydroxide) in the transcript. Note: NH₄OH is typically classified as a base (ammonium hydroxide solution) rather than a salt in standard chemistry, but it appears in the transcript as an example under salts.

  • May contain polyatomic ions; dissolution behavior in water depends on ionic nature.

• Molecules

  • Covalent compounds; contain only nonmetals.

  • Examples: Br₂ (bromine, a diatomic molecule) and CH₄ (methane).

• Acids

  • Acids begin with an H in their formula (as stated in the transcript).

  • Example provided: HCl (hydrochloric acid).

  • (Additional standard context beyond the transcript: in aqueous solution, acids increase H⁺ concentration, often written as H₃O⁺ or simply H⁺; they are typically characterized by proton donation.)

• Bases

  • Bases end with OH in their formula (as stated in the transcript).

  • Examples: NaOH (sodium hydroxide) and Ca(OH)₂ (calcium hydroxide).

  • (Standard context: bases donate OH⁻ in solution or accept H⁺; in aqueous solution they raise OH⁻ concentration.)

• Quick classification exercise from the transcript

  • CuCl₂ → Salt (ionic compound).

  • HC₂H₃O₂ (acetic acid) → Acid.

  • Ca(OH)₂ → Base.

  • Note: The transcript also mentions dissolution in water and aqueous solutions (see solubility section) and provides a format for labeling substances as salt, molecule, acid, or base.

• Dissolves in water; aqueous notation

  • Dissolving in water yields an aqueous solution, denoted as (aq).

  • The transcript uses the terms: soluble (dissolves in water) vs insoluble (does not dissolve in water; solid, S).

  • The classification exercise includes noting when a substance dissolves in water (soluble) and forms an aqueous solution.

Solubility in water: soluble vs insoluble

• Soluble vs Insoluble

  • Soluble: dissolves in water; forms an aqueous solution (aq).

  • Insoluble: does not dissolve in water; remains a solid (S).

• Ionic compounds and water dissolution rules (as stated in the transcript)

  • Ionic compounds will dissolve in water if they contain:

  • An alkali metal ion (group 1) (e.g., Li⁺, Na⁺, K⁺, etc.).

  • Ammonium ion (NH₄⁺).

  • Nitrate ion (NO₃⁻).

  • The transcript hints (with partial examples) that CO₃²⁻ (carbonate) might be included in solubility considerations in some contexts, though the clearest rule cited is about group 1 metals, NH₄⁺, and NO₃⁻.

  • NO₃⁻ (nitrate) is explicitly listed as a soluble ion in the transcript.

  • CO₃²⁻ (carbonate) is mentioned in passing in the transcript (likely as an example of another polyatomic ion to consider), but standard solubility rules state that carbonate salts are generally insoluble except with group 1 cations and NH₄⁺.

• Other ionic substances

  • The transcript notes that there are additional ionic substances that dissolve in water, but students are not expected to memorize all of them; a solubility chart can be used as a reference.

• General remarks about covalent substances

  • The transcript indicates that most covalent substances do not dissolve in water.

  • It also notes that rules for covalent solubility will be learned later in the course.

• Specific example mentions in the transcript

  • Ca(OH)₂ is cited as a base (which, in standard chemistry, is an ionic hydroxide that can be soluble to varying extents depending on the specific base).

  • The discussion emphasizes that water solubility is a key criterion for labeling substances as soluble (aq) vs solid (S).

• Real-world relevance and connections

  • Understanding solubility helps predict reaction outcomes in aqueous environments, including precipitation reactions, acid-base interactions, and nutrient uptake in biological systems.

  • Solubility rules form a practical heuristic in qualitative inorganic analysis and in planning synthesis and purification steps in chemistry labs.

• Caveats and learning trajectory

  • While the transcript provides a simplified set of rules for ionic solubility (group 1 cations, NH₄⁺, and NO₃⁻ as soluble), students should learn the complete solubility chart in later coursework, which includes exceptions (e.g., common ion pairs, sulfate solubility exceptions, etc.).

  • Covalent solubility is more nuanced and context-dependent (e.g., very polar covalent substances like alcohols and sugars may dissolve, while many nonpolar covalent substances do not).

• Summary takeaways

  • Classification categories: salts (ionic), molecules (covalent), acids, bases.

  • Acids start with H; bases end with OH.

  • Solubility in water is a practical consideration for predicting whether a substance will be aqueous (aq) or solid (S).

  • Aqueous dissolution of ionic compounds often depends on the presence of alkali metals (group 1) ions, ammonium NH₄⁺, and nitrate NO₃⁻.

  • The transcript emphasizes using a solubility chart for broader rules and notes that covalent solubility rules will be introduced later.