Co←mmpound Naming: Ionic, Covalent, and Acids O

Ionic Compounds Naming

General idea: ionic compounds are formed from a metal (positive ion) and a non-metal (negative ion).
Naming rules:
- Metal comes first.
- Non-metal name ends with "-ide" when written as the anion (e.g., sulfur → sulfide, hydrogen → hydride).
- The formula reveals the subscripts (how many atoms of each element), not just the name.
Writing formulas from names (and charges):
- Find the charges of the ions.
- Some charges are easy to recall from the periodic table: alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17).
- Other ions: chalcogens (Group 16), pnictogens (Group 15), crystallogens (older term; Group 14 and nearby in early periods).
- For less obvious ions, consult the periodic table’s charge hints on the back, or the common-ion sheet.
- Swap-and-drop (criss-cross) method to balance charges, then simplify to smallest whole-number ratio.
Key concept: electrical neutrality of ionic compounds.
Common shorthand: write the empirical formula that reflects the smallest whole-number ratio of ions.
Examples of process:
- If one metal cation with a 1+ charge pairs with one oxide anion with a 2− charge: balancing gives a 1:1 ratio? (example: NaCl). In general, use the criss-cross method: for Na^+ and Cl^−, formula is $ext{NaCl}$ .
- For Ca^{2+} and NO3^{−} (nitrate), balancing gives Ca(NO3)2: the polyatomic ion NO3 is treated as a unit, bracketed when needed (see polyatomic ions section).
Polyatomic ions and bracket notation (when the anion is polyatomic):
- If the anion ends in -ate or -ite, it’s a polyatomic ion (e.g., NO3^−, SO4^{2−}, CO_3^{2−}).
- Treat the whole polyatomic ion as a single unit when balancing.
- If balancing requires a -2 or another charge, brackets are used around the polyatomic ion with a subscript outside the bracket.
- Example: Calcium nitrate → $ext{Ca(NO}3 ext{)}2$ .
Ammonium is a positive polyatomic ion: $ext{NH}_4^+$ .
Multivalent ions (variable oxidation states):
- Many transition metals (and some heavier elements) have more than one common oxidation state.
- Check the common-ion sheet for their charges; some exceptions exist (e.g., Silver is typically +1).
- Zinc and other Group 12 elements are not considered transition metals for these purposes.
- When naming multivalent metal ions, specify the charge with Roman numerals: e.g., Copper (II) chloride (CuCl_2) vs. Copper (I) chloride (CuCl).
Practice examples (formula → name):
- $ext{NaCl} ightarrow ext{Sodium chloride}$
- $ext{K}_2 ext{O} ightarrow ext{Potassium oxide}$
- $ext{Be}3 ext{N}2 ightarrow ext{Beryllium nitride}$
- $ext{AlH}_3 ightarrow ext{Aluminum hydride}$
- $ext{SrO} ightarrow ext{Strontium oxide}$
- $ext{AgF} ightarrow ext{Silver fluoride}$
- $ext{ZnCl}_2 ightarrow ext{Zinc chloride}$
Practice problems (names you’d write from the formulas):
- $ext{NaCl} ightarrow ext{Sodium chloride}$
- $ext{K}_2 ext{O} ightarrow ext{Potassium oxide}$
- $ext{Be}3 ext{N}2 ightarrow ext{Beryllium nitride}$
- $ext{AlH}_3 ightarrow ext{Aluminum hydride}$
- $ext{SrO} ightarrow ext{Strontium oxide}$
- $ext{AgF} ightarrow ext{Silver fluoride}$
- $ext{ZnCl}_2 ightarrow ext{Zinc chloride}$

Polyatomic Ions and Multivalents (Page 5 details)

Polyatomic ions:
- If the anion ends in -ate or -ite, treat it as a single unit; it has a charge (commonly -1, -2, etc.).
- When balancing, bracket the polyatomic ion if you need to place a numerical subscript outside the bracket.
- Example: $ext{Ca(NO}3 ext{)}2$ (calcium nitrate) is formed from $ext{Ca}^{2+}$ and $ext{NO}_3^{-}$ .
Ammonium as a cation: ammonium is a polyatomic cation, $ext{NH}_4^+$ .
Multivalent ions:
- Many transition metals show multiple oxidation states.
- Use common-ion sheets to determine possible charges.
- Examples:
- Silver commonly +1: $ext{Ag}^+$ .
- Zinc (and Group 12 elements) typically have a fixed oxidation state and are not treated as multivalent in this context.
- Naming multivalent ions uses Roman numerals: e.g., Copper (II) chloride, $ext{CuCl}_2$ .

Practice: Ionic Formulas and Names (Page 6)

Given formulas and names (interpretations):
- $ext{AgClO}$ → Silver hypochlorite
- $ext{K}2 ext{SO}4$ → Potassium sulfate
- $ext{CuCl}_2$ → Copper(II) chloride
- $ext{NH}_4 ext{Cl}$ → Ammonium chloride
- $ext{Cu}3( ext{PO}4)_2$ → Copper(II) phosphate
- $ext{Fe(HCO}3)3$ → Iron(III) bicarbonate
- $ext{(NH}4) ext{C}2 ext{H}3 ext{O}2$ → Ammonium acetate
- $ext{U(CN)}_4$ → Uranium(IV) cyanide

Covalent Compounds Naming (Page 7)

Covalent compounds form from non-metals.
Naming rules:
- Non-metals first? In many cases, the more metallic/less electronegative element is named first, but the general rule is non-metal/non-metal exchange and prefix usage.
- The last non-metal changes ending to "-ide" (e.g., sulfur → sulfide, hydrogen → hydride).
- Use prefixes to indicate the number of each element (mono-, di-, tri-, etc.).
- Do not use a prefix for the first element if there is only one of that element.
Writing formulas from names (for covalent compounds):
- Unlike ionic names, the covalent names indicate the exact number of each element.

Covalent Formulas and Names Practice (Page 8)

Examples (formulas):
- $ext{SO}_2 ightarrow ext{Sulfur dioxide}$
- $ext{CO} ightarrow ext{Carbon monoxide}$
- $ext{NO}_2 ightarrow ext{Nitrogen dioxide}$
- $ext{CO}_2 ightarrow ext{Carbon dioxide}$
- $ext{SF}_6 ightarrow ext{Sulfur hexafluoride}$
- $ext{XeF}_8 ightarrow ext{Xenon octafluoride}$
- $ext{B}2 ext{H}4 ightarrow ext{Diborane}$

Covalent Names Practice (Page 9)

Given names, the formulas (or vice versa):
- Sulfur tetrachloride → $ext{SCl}_4$
- Trihydrogen carbon chloride → ambiguous in the slide; likely CH3Cl (chloromethane) if referring to common covalent naming; formula: $ext{CH}3 ext{Cl}$
- Carbon monoxide → $ext{CO}$
- Phosphorus trihydride → $ext{PH}_3$
- Iodine heptafluoride → $ext{IF}_7$
- Chlorine monofluoride → $ext{ClF}$

Exceptions and Common Names (Page 10)

Common exceptional covalent names:
- Water = $ext{H}_2 ext{O}$
- Methane = $ext{CH}_4$
- Ammonia = $ext{NH}_3$
Note: These are widely used traditional/common names beyond the straightforward naming rules.

Naming Acids (Page 11)

All acids contain hydrogen; recognizing acids is easy when a hydrogen is present.
Binary acids (hydrogen + a halogen or hydrogen + another element): H–A → hydro- A-ic acid.
- Example: $ext{HCl} ightarrow ext{hydrochloric acid}$
Oxyacids (hydrogen with a polyatomic anion): H–polyatomic ion.
- If the polyatomic anion ends in -ate, the acid name ends in -ic acid.
- If the polyatomic anion ends in -ite, the acid name ends in -ous acid.
- Examples:
- $ext{H}2 ext{SO}4 ightarrow ext{sulfuric acid}$
- $ext{HNO}_2 ightarrow ext{nitrous acid}$
Quick recap for acids:
- H + element → hydro- [element stem]-ic acid (binary acids).
- H + polyatomic -ate → [polyatomic stem]-ic acid.
- H + polyatomic -ite → [polyatomic stem]-ous acid.

Quick Reference: Key Formulas (LaTeX)

Ionic compounds (selected examples):
- $ext{NaCl}$ , $ext{K}2 ext{O}$ , $ext{Be}3 ext{N}2$ , $ext{AlH}3$ , $ext{SrO}$ , $ext{AgF}$ , $ext{ZnCl}_2$
- Nomenclature rules: metal first, non-metal with -ide; brackets + subscripts for polyatomic ions when needed; Roman numerals for multivalents.
Polyatomic ions and bracket notation:
- $ext{Ca(NO}3 ext{)}2$ (calcium nitrate)
- Ammonium ion: $ext{NH}_4^+$
Covalent compounds (selected):
- $ext{SO}2$ , $ext{CO}$ , $ext{NO}2$ , $ext{CO}2$ , $ext{SF}6$ , $ext{XeF}8$ , $ext{B}2 ext{H}_4$
Covalent names:
- Sulfur tetrachloride → $ext{SCl}_4$
- Phosphorus trihydride → $ext{PH}_3$
- Iodine heptafluoride → $ext{IF}_7$
- Chlorine monofluoride → $ext{ClF}$
Acids (examples):
- $ext{HCl} ightarrow ext{hydrochloric acid}$
- $ext{H}2 ext{SO}4 ightarrow ext{sulfuric acid}$
- $ext{HNO}_2 ightarrow ext{nitrous acid}$
Special cases: water, methane, ammonia formulae: $ext{H}2 ext{O}, ext{CH}4, ext{NH}_3$

Title: Compound Naming: Ionic, Covalent, and Acids (Ionic/Covalent/Naming A-chem topics)