Bonding, Nomenclature, and Polyatomic Ions — Lecture Notes

Fixed oxidation states and common charges (as discussed in lecture)

Zinc (Zn) is always +2; it loses two electrons.
Cadmium (Cd) is also +2.
Silver (Ag) is +1; it loses one electron.
Cobalt (Co) and Nickel (Ni) commonly show +2 (Co^{2+}, Ni^{2+}).
Manganese is often listed as +2, but there are multiple oxidation states; +2 is the common case, with other states possible.
Copper (Cu) can be +1 or +2; examples discussed include copper(I) and copper(II).
Iron (Fe) can be +2 or +3; Fe^{2+} or Fe^{3+} discussed; iron(III) is common in oxides.
Gold (Au) can be +1 or +3.
Chromium (Cr) can be +3 or +6 (varies by compound).
Chlorine (Cl) was noted as +2 in this context by the instructor (context-specific; standard oxidation states of Cl are usually -1, but the guide states +2 here).
Platinum (Pt) is +2 in the exampleimet.
Aluminum (Al) is +3.
Gallium (Ga) and Indium (In) are +3 in the examples provided.
A “six-pack” of commonly encountered fixed-oxidation-state species mentioned: Ag, Zn, Cd, Al, Ga, In (and notes that these elements typically display a common oxidation state in many compounds).
Some elements have variable oxidation states (e.g., copper, iron, chromium) and we’ll make best-guess determinations with +3, +4, etc., when needed.
Practical takeaway: these fixed charges help predict formulas and naming, especially for ionic compounds and for balancing charges.

The stair-step and the metals/nonmetals distinction

The staircase in the periodic table separates metals from nonmetals; this distinction helps predict bonding type and naming conventions.
Metals lie to the left of the staircase; nonmetals to the right.
This separation underpins the distinction between ionic (metal + nonmetal) and covalent (two nonmetals) bonding.

Polyatomic ions and pattern recognition (oxyanions)

Polyatomic ions are groups of atoms that act as a single unit with a net charge; they are not split apart when forming compounds.
Common polyatomic ions highlighted (bold items on the sheet):
- Ammonium: \NH{4}^{+}\n - Hydroxide: \OH^{-}\n - Nitrate: \NO{3}^{-}\n - Nitrite: \NO{2}^{-}\n - Sulfate: \SO{4}^{2-}\n - Sulfite: \SO{3}^{2-}\n - Phosphate: \PO{4}^{3-}\n - Perchlorate: \ClO{4}^{-}\n - Chlorate: \ClO{3}^{-}\n - Carbonate: \CO_{3}^{2-}\n - (Other common polyatomics mentioned include bicarbonate, acetate, ammonium, etc.)
Pattern for oxyanions (the “ate” vs “ite” rule):
- When you reduce the oxygen by one, you typically get the “ite” form with the same overall charge.
- Example: \NO{3}^{-} (nitrate) vs. \NO{2}^{-} (nitrite); \SO{4}^{2-} (sulfate) vs. \SO{3}^{2-} (sulfite).
Prefixes for oxyanions: perchlorate (ClO4^-), chlorate (ClO3^-), chlorite (ClO2^-), hypochlorite (ClO^-).
For acids derived from oxyanions, add the word “acid” and adjust the suffix:
- Ate → ic acid (e.g., sulfate → sulfuric acid)
- Ite → ous acid (e.g., sulfite → sulfurous acid)
- Example acids: \H{2}SO{4} is sulfuric acid; \H{2}SO{3} is sulfurous acid; \H{3}PO{4} is phosphoric acid; \HNO{3} is nitric acid; \HNO{2} is nitrous acid.
In salts with a polyatomic anion, the polyatomic ion is not altered (no -ide on the polyatomic portion); the metal portion is what changes to balance charge.
- Example: sodium bicarbonate = \NaHCO{3} (ammonium bicarbonate would be NH{4}HCO{3}); calcium sulfate = \CaSO{4}; aluminum carbonate = \Al{2}(CO{3})_{3} (no modification to carbonate).
The ions listed above are very helpful for quick naming and formula-building; memorization is helpful but the pattern and the sheet are the intended tools.

Ionic vs Covalent bonding and the naming implications

Ionic bonds: formed by transfer of electrons from a metal to a nonmetal; electrostatic attraction between cation and anion; high melting points, high boiling points, crystalline solids (e.g., table salt).
Covalent bonds: formed when two nonmetals share electrons; lower melting/boiling points; often non-crystalline or molecular in nature.
Transfer vs sharing rules (basic view):
- Metal + Nonmetal (ionic): transfer of electrons; typically forms salts with fixed charges that balance to neutrality.
- Two nonmetals (covalent): sharing of electrons; electrons are not fully transferred.
Exceptions and nuance exist; hydrogen can act like a metal or a nonmetal depending on context, which can influence naming and classification in borderline cases.
Example discussions:
- Sodium (Na, metal) + Chlorine (Cl, nonmetal) → ionic compound NaCl (sodium chloride).
- Calcium (Ca, metal) + Oxygen (O, nonmetal) → typically ionic CaO (calcium oxide). If both partners are nonmetals, e.g., Cl and O in a molecule, you’re in covalent territory.
- Calcium carbonate CaCO_{3} as an ionic compound containing a polyatomic anion (carbonate) with a metal.
Covalent naming tends to reflect the actual sharing and may use prefixes (mono-, di-, tri-, etc.) for two-nonmetal compounds; ionic naming uses the metal name (and oxidation state if needed) plus the nonmetal name with -ide, or the polyatomic ion name.

How to name binary and polyatomic compounds (nomenclature rules)

Binary compounds are those with only two different elements. They can be:
- Binary ionic (metal + nonmetal): name the metal first, then the nonmetal with -ide suffix (e.g., sodium iodide, NaI).
- Binary molecular (two nonmetals): use prefixes to denote the number of atoms (e.g., CO to carbon monoxide, CO_{2} to carbon dioxide).
If a binary compound contains a metal with variable oxidation state, a Roman numeral is used to indicate the charge (e.g., CuCl is copper(I) chloride; CuCl_{2} is copper(II) chloride).
For polyatomic compounds (more than two different elements, or a metal + a polyatomic ion):
- Do not alter the polyatomic ion; name the metal first (if present) and then name the polyatomic ion as is (e.g., calcium carbonate, CaCO{3}; ammonium chloride, NH{4}Cl).
- If the formula includes a polyatomic ion, you typically do not apply the “-ide” ending to the polyatomic portion.
Example naming walk-throughs from the lecture:
- Potassium acetate: \C{2}H{3}O_{2}^{-} with a K^{+} result; name: potassium acetate; the acetate ion is a polyatomic ion, not modified to -ide.
- Ammonium chloride: NH_{4}^{+} with Cl^{-}; name: ammonium chloride (not a hydrogen chloride naming for the chloride portion).
- Sodium bicarbonate: NaHCO_{3}; bicarbonate is the polyatomic ion; name: sodium bicarbonate.
- Calcium sulfate: CaSO_{4}; name: calcium sulfate.
- Aluminum carbonate: Al{2}(CO{3})_{3}; carbonate is the polyatomic ion; name: aluminum carbonate.
Recognizing a binary ionic vs a binary molecular compound can be aided by inspecting the name or the formula:
- If the name contains a Roman numeral, or the formula requires balancing to achieve neutrality using typical ion charges, it’s ionic (often binary, sometimes with a polyatomic anion).
- If the name uses prefixes (mono-, di-, tri-) to indicate two nonmetals, it is typically covalent (molecular).
Special note: not all endings imply ionic vs covalent; for example, endings like -ide can occur in both contexts depending on whether a polyatomic ion is involved or whether the bond is ionic or covalent.

Acids and how they are named from ions

Acids come in two broad categories:
- Polyatomic-acid derived from oxyanions (e.g., nitrate, sulfate, phosphate): the base anion name changes suffix to indicate acid form, and the word “acid” is appended.
- Binary acids (hydrogen paired with a nonmetal, e.g., HCl, H2S): named with the hydro- prefix, the root of the nonmetal, and -ic or -ous as appropriate, followed by acid.
Polyatomic acids naming pattern (for oxyanions):
- For anions ending in -ate, the corresponding acid ends with -ic (e.g., sulfate -> sulfuric acid).
- For anions ending in -ite, the corresponding acid ends with -ous (e.g., sulfite -> sulfurous acid).
- Examples:
- NO{3}^{-} (nitrate) → HNO{3} (nitric acid).
- NO{2}^{-} (nitrite) → HNO{2} (nitrous acid).
- SO{4}^{2-} (sulfate) → H{2}SO_{4} (sulfuric acid).
- SO{3}^{2-} (sulfite) → H{2}SO_{3} (sulfurous acid).
- PO{4}^{3-} (phosphate) → H{3}PO_{4} ( phosphoric acid).
- ClO{4}^{-} (perchlorate) → HClO{4} (perchloric acid).
- ClO{3}^{-} (chlorate) → HClO{3} (chloric acid).
- ClO^{-} (hypochlorite) → HClO (hypochlorous acid).
Binary acids naming pattern (hydrogen with a nonmetal):
- Hydro + root of the nonmetal + -ic + acid (e.g., HCl is hydrochloric acid; HBr is hydrobromic acid).
- If the anion is derived from hydrogen sulfide (H_{2}S), the acid is hydrosulfuric acid (older nomenclature) or simply hydrogen sulfide in pure form (used in gas context); in common binary acid naming, the hydro-prefix pattern is used.
The instructor emphasized the following practical rules when naming acids:
- If the compound contains hydrogen and a polyatomic anion (oxyanion), use the polyatomic acid naming pattern described above (ate → -ic, ite → -ous).
- If the compound is a binary acid, use the hydro- naming convention.

Using the periodic table to predict bonding type and formulate compounds

Example workflow from lecture:
- Given Na and Cl, determine whether the compound is ionic or covalent by metal vs nonmetal placement on the periodic table (Na is a metal; Cl is a nonmetal) → ionic compound (NaCl).
- Use the charges to balance and determine formula: Na^+ (group 1) and Cl^- (group 17) balance 1:1 → NaCl.
- For Ca and O: both are shown; Ca is a metal (left side) and O is nonmetal; ionic CaO; Ca^2+ and O^{2-} balance 1:1.
Deductive method (the “plate detector” method) when given a formula but not the oxidation states:
- Oxygen is almost always -2.
- If a compound contains O and a metal, use the metal’s oxidation state to balance charges with O^{2-} to achieve neutrality.
- Example: Cu{2}O. Oxygen is -2; to balance, two Cu must supply +2 total; each Cu is +1 → copper(I) oxide, Cu{2}O.
- Example: CuCl{2}. Chlorine is -1; to cancel, copper must be +2 → copper(II) chloride, CuCl{2}.
- Example: Fe{2}O{3}. Oxygen is -2; total -6 from O; to balance, Fe must contribute +3 total; two Fe give +6 → iron(III) oxide, Fe{2}O{3}.
Practical caveats mentioned in lecture:
- Mercury compounds can be tricky (Mercury(I) Hg_{2}^{2+} or Mercury(II) Hg^{2+}); the instructor warned these can be confusing and are rarely asked on the exam.
- Tin (Sn) also has multiple oxidation states (e.g., Sn^{2+}, Sn^{4+}); similar balancing rules apply.
- Some compounds with multiple elements may be polyatomic (e.g., CaSO{4}, CaCO{3}, Ca{3}(PO{4})_{2}); the polyatomic group is not broken up when balancing.
Everyday ionic examples mentioned:
- Antacids, bleach, toothpaste frequently involve ionic compounds.
- The discussion referenced real-world cases like environmental issues (e.g., Erin Brockovich) to illustrate the impact of chemical compounds and environmental health concerns.

Key formulas, constants, and study aids mentioned (to be used on the exam)

Density (example formula):
ho = rac{m}{V}
Manipulation practice: if given mass and density, find volume: V = rac{m}{
ho}
Gas volume at STP: $V = 22.4\ ext{L/mol}$ per mole of an ideal gas (gas laws context).
Temperature conversion (general knowledge for gas problems): $^ ext{°C} = \frac{5}{9} (^ ext{°F} - 32)$
Back-of-book formulas (by chapter progression): density; percent mass; percent yield; various gas-law equations; heat capacity relationships (to be covered in later chapters).

Polyatomic ions and memorization tips

The following polyatomic ions are especially helpful to memorize or recognize quickly:
- Ammonium: \NH{4}^{+}\n - Hydroxide: \OH^{-}\n - Nitrate: \NO{3}^{-}\n - Nitrite: \NO{2}^{-}\n - Sulfate: \SO{4}^{2-}\n - Sulfite: \SO{3}^{2-}\n - Phosphate: \PO{4}^{3-}\n - Carbonate: \CO{3}^{2-}\n - Acetate: \C{2}H{3}O{2}^{-} (an example of a conjugate polyatomic ion that appears in salts)
Recognize that the polyatomic ions stay intact as units when forming compounds and that the metal portion is the part that changes to balance charges.

Study and course context from the lecture

The instructor provided on-board sheets with polyatomic ions, metric prefixes, constants, and some conversions for use during the exam.
Not all prefixes need to be memorized, but know the basic ones and how to apply them when naming compounds.
The front sheet also includes constants (like Avogadro’s number context) and common conversions for gas and heat problems.
The lecture emphasizes that you will be using algebra in stepwise ways to manipulate formulas, especially with density and gas-law relationships.
The instructor mentions that the next topics will focus on polyatomic ions, binary acids, and hydrates, with Chapter 3 (math-heavy) and Chapter 4 (continued math focus) being central to understanding.
Exam scope: Chapter 1 and Chapter 2 concepts are in the near-term exam; Chapter 3 (and possibly some 3.1 Section material) may appear; prelab due Monday; a focus on the math-heavy content in Chapters 3–5 is anticipated.
Real-world context: the instructor connects chemical knowledge to environmental health issues (e.g., water contamination and legal cases) to illustrate the broader importance of chemistry knowledge.

Quick reference reminders (cheat-sheet style)

Ionic vs Covalent: Metals + nonmetals (ionic) vs two nonmetals (covalent).
Balancing charges in ionic compounds often uses crisscross of charges to determine subscript numbers (e.g., Na^+ and Cl^- → NaCl; Fe^{3+} and O^{2-} → Fe{2}O{3}).
For acids, use -ic/-ous endings for polyatomic oxyanions; use hydro- prefix for binary acids.
For polyatomic ions, do not break apart the group when naming; treat as a unit.
Use the stair-step to anticipate bonding type and naming conventions.
Density and gas-law related math will be essential in subsequent chapters; practice algebraic manipulation.