Nomenclature and Bonding Concepts from

Overview and Exam Approach

• The session focuses on nomenclature and the quiz content; the instructor emphasizes that naming compounds is mainly memorization, so spread memorization over a week rather than cramming two days before the test.
• Learning approach highlighted: You can practice a lot with worksheets; the key is to show your thinking process on problems rather than just providing the final answer.
• Grading tip shared: For a problem worth 5 points, writing the setup and reasoning (electrons, neutrons, protons, etc.) can earn most of the points; the final answer itself contributes relatively little compared to showing the reasoning.
• Example of signal for correct credit: When solving for significant figures and other details, the method matters as much as the final numeric result.
• Tools allowed: For exams (online or in-class), calculators are allowed; the instructor mentions that you can use a calculator without penalty.
• Textbook/notes hint: PowerPoints and “iPods” (likely digital notes) are provided; PowerPoints are useful for memorization but not ideal for deep learning; the emphasis is on memory-based review, with the expectation that students will memorize the rules ahead of time.

Periodic Table Trends and Hydrogen

• Families on the periodic table share similar properties; a central region of the table is often disregarded for simplifying rules during the lecture.
• Hydrogen (H) is unusual: most of the time it behaves as if it were in the same column as the alkali metals (Group 1), but occasionally it behaves like a halogen (Group 17). Overall, it’s treated as acting like a nonmetal for practical purposes, though it’s a special case.
• The alkali metals (Group 1 elements) are described as solid metals that react vigorously with water to produce hydrogen gas; this is a characteristic property demonstrated early in the course.
• The lecturer references a “stair-step” (the dividing line between metals and nonmetals) and notes metalloids lie around this boundary; hydrogen sits near or above this boundary in discussions.
• The speaker notes some confusion around certain elements (e.g., aluminum) being labeled as metals vs. metalloids in the slide, highlighting that there can be nuanced classifications in early materials.
• The rest of the elements left of the staircase are metals; those to the right are nonmetals; metalloids sit along the staircase.

Covalent vs Ionic Bonding and Bonding Trends

• Bond type depends on the elements involved:
  • Covalent bonds form mainly between nonmetals (e.g., C and O in CO or CO₂).
  • Ionic bonds form from a metal and a nonmetal (e.g., NaCl, NaNO₂), where electrons are transferred to form ions.
• The initial step in deciding the type of compound is to identify whether the constituent elements are metals or nonmetals:
  • If a metal combines with a nonmetal, the compound tends to be ionic.
  • If both are nonmetals, the compound tends to be covalent.
• Example discussed: Na and Cl form an ionic compound; the goal is to decide ionic vs covalent first, then apply naming conventions.
• Group and block ideas mentioned (though sometimes imprecise in the lecture):
  • The text references s, p, and other blocks and suggests a practical approach based on the position of elements on the periodic table.
  • Transition metals (the d-block) often form multiple oxidation states and are common in some compounds; main-group elements (left and center) tend to have more predictable oxidation states.
• Covalent bond example: CO and H₂O involve nonmetals; the bond is covalent because the bonded atoms are nonmetals.
• The lecture emphasizes that many of the compounds students name in chemistry (especially simple binary compounds) have covalent bonds when nonmetals are involved, and ionic bonds when a metal is involved.

Ions, Polyatomic Ions, and Oxidation States

• An ion is formed by the loss or gain of electrons; single atoms can form monatomic ions, while groups of atoms can form polyatomic ions (a group of atoms with an overall charge acting as a single unit).
• Polyatomic ions example:
  • Phosphate: ext{PO}_4^{3-} (a polyatomic ion with a 3− charge). The group acts as a single unit in compounds, and the charge is assigned to the ion as a whole, not to individual atoms.
  • Acetate: ext{CH}3 ext{COO}^- ext{ or } ext{C}2 ext{H}3 ext{O}2^- (common polyatomic ion used in naming; highlighted as more relevant in the second semester).
• The charges on polyatomic ions come from the overall loss or gain of electrons by the group; in some cases, it’s not clear which exact atom contributed the electrons—the ion is treated as a single unit with a net charge.
• Some charge values for main-group elements (oxidation states) are discussed:
  • Group 1: typically +1
  • Group 2: typically +2
  • Group 3: typically +3
  • Group 4: can be +4 or −4 (less favorable to form ions, but possible for some elements such as Si and Ge at the top portion of the group)
  • Group 5: typically −3
  • Group 6: typically −2
  • Group 7: typically −1
  • Group 8: typically 0 (noble gases, usually don’t form ions)
• A note from the lecturer: the “top” of Group 4 (e.g., Si, Ge) can exhibit +4 oxidation states, while bottom portions may exhibit different tendencies; this illustrates why predicting charges for some elements can be tricky.
• The speaker cautions that many of these ion charge patterns are best memorized for the exam, but there are underlying reasons tied to electron configurations and valence electrons.

Rules of Nomenclature and the -ide Concept

• Naming approach for ionic compounds typically starts by deciding ionic vs covalent bonding, then applying a fixed set of rules for naming:
  • For binary ionic compounds (a metal and a nonmetal), the metal name is followed by the nonmetal name with the suffix -ide on the nonmetal: e.g., sodium chloride (NaCl), where chloride is the nonmetal with -ide suffix.
  • For polyatomic ions in salts, the name uses the polyatomic ion name (e.g., sodium nitrate NaNO₃, where nitrate is NO₃⁻).
• The suffix -ide is associated with the anion (nonmetal or polyatomic ion’s ending is typically -ide for simple binary ions; more complex ions retain their certain suffixes like -ate or -ite when applicable).
• The lecture mentions that the -ide ending corresponds to the single nonmetal component in a binary compound; more complex naming (with polyatomic ions) follows standard ion names.
• Example mentioned: Sodium nitrite (NaNO₂) uses the nitrite polyatomic ion (NO₂⁻).
• The instructor notes that there is a widely used list of polyatomic ions (often provided on a cheat sheet or on the periodic table handout) that exams reference; using ions outside that list may result in lost points on some questions.

Hydrates and Dot Notation

• Hydrates are ionic compounds that have water molecules incorporated into their crystal structure; the water is not covalently bonded to the ions but is part of the crystal lattice (often labeled as water of crystallization).
• Notation: the water molecules are indicated with a dot, e.g.,
  • Cobalt chloride hexahydrate: ext{CoCl}2 \cdot\ 6 ext{H}2 ext{O}
  • In general, a hydrate is written as: formula
  • Example: ext{CuSO}4 \, ext{(dot)} \, 5 ext{H}2 ext{O} for copper(II) sulfate pentahydrate.
• The speaker uses the term “hexahydrin” to indicate six water molecules associated with cobalt, illustrating the concept of hydrates, though the actual common hydrate form might be referred to as cobalt chloride hexahydrate ( ext{CoCl2·6H2O}) in many texts.
• The dot notation helps distinguish the crystalline water from the primary formula unit of the salt.

Examples and Practical Notes Discussed in the Transcript

• Example of composition and naming:
  • Sodium chloride: a classic ionic salt formed from Na⁺ and Cl⁻; name uses the metal (sodium) + nonmetal with -ide (chloride).
  • Sodium nitrite: NaNO₂; contains the nitrite polyatomic ion NO₂⁻; name reflects the polyatomic ion as-is.
  • Phosphate ion: ext{PO}_4^{3-}; common in salts such as sodium phosphate, calcium phosphate, etc.
  • Acetate ion: ext{CH}3 ext{COO}^- or ext{C}2 ext{H}3 ext{O}2^-; used in various salts and esters; its naming can appear in more advanced topics (noted as more common in second semester).
  • Aluminum fluoride: ext{AlF}_3; a typical ionic compound formed from a +3 metal cation with a fluoride anion.
• The instructor notes that many questions in ACS exams involve recognizing polyatomic ions and applying standard naming conventions. The correct approach is to check a provided list of allowed ions on the exam or in the study materials.
• There is an emphasis on recognizing that some bonds within polyatomic ions are covalent, even though the overall compound is ionic, because the ions themselves are composed of covalently bonded atoms that share electrons within the ion.
• A reminder that certain elements (especially around the boundary between metals and nonmetals) can exhibit tricky behavior in naming and bonding; metalloids sit near the dividing line and may show mixed characteristics.

Quick Reference: Key Formulas and Names (LaTeX)

• Polyatomic ions:
  • Phosphate: ext{PO}_4^{3-}
  • Nitrate: ext{NO}_3^{-}
  • Nitrite: ext{NO}_2^{-}
  • Acetate: ext{CH}3 ext{COO}^- or ext{C}2 ext{H}3 ext{O}2^-
  • Sulfate: ext{SO}_4^{2-}
• Example salts:
  • Sodium chloride: ext{NaCl}
  • Sodium nitrite: ext{NaNO}_2
  • Sodium phosphate: ext{Na}3 ext{PO}4
• Hydrates (dot notation):
  • Cobalt chloride hexahydrate: ext{CoCl}2 \cdot ext{6H}2 ext{O}
  • Copper(II) sulfate pentahydrate: ext{CuSO}4 \, ext{·}\,5 ext{H}2 ext{O}

Practice Tips and Study Recommendations

• Use the provided worksheets to practice naming and recognize patterns; memorization should be spaced out across at least a week.
• For online exams, you may use a calculator as needed, which can help with arithmetic accuracy but does not replace understanding the rules.
• Focus on the rule-set for naming: determine covalent vs ionic first, then apply the appropriate naming convention (ide suffix for simple binary anions, polyatomic ion names for salts).
• Build a solid mental map of common polyatomic ions (e.g., nitrate, nitrite, sulfate, phosphate, acetate) to reduce errors during naming and formula writing.
• Review the periodic table’s left-right division: metals on the left (tend to form cations when possible), nonmetals on the right (tend to form anions or covalently bonded species); metalloids sit along the staircase and can show mixed behavior.
• Remember that hydrogen’s behavior is context-dependent; it can act like a nonmetal most of the time, but may resemble a Group 1 element in some contexts.
• When unsure about a species, check if it fits a common pattern (binary ionic compound with a metal and a nonmetal vs a covalently bonded molecule of two nonmetals), and use the standard naming rules accordingly.
• Always reference the allowed ions list or the standard polyatomic ion table provided with the course materials to ensure correct naming for exam problems.