Common Ions, Nomenclature, and Acid Naming – AP Chemistry Study Notes

Common Ions to Memorize

• Success in AP Chemistry presumes instantaneous recall of every ion, its formula, and charge on Day 1.
• A periodic table will always be supplied; positions of elements make many charges “automatic.”

Monatomic Cations (fixed charge)

• $\mathrm{H^+}\;\text{hydrogen}$
• $\mathrm{Li^+}\;\text{lithium}$
• $\mathrm{Na^+}\;\text{sodium}$
• $\mathrm{K^+}\;\text{potassium}$
• $\mathrm{Rb^+}\;\text{rubidium}$
• $\mathrm{Cs^+}\;\text{cesium}$
• $\mathrm{Be^{2+}}\;\text{beryllium}$
• $\mathrm{Mg^{2+}}\;\text{magnesium}$
• $\mathrm{Ca^{2+}}\;\text{calcium}$
• $\mathrm{Sr^{2+}}\;\text{strontium}$
• $\mathrm{Ba^{2+}}\;\text{barium}$
• $\mathrm{Al^{3+}}\;\text{aluminum}$
• Transition/ d-block cations with fixed charge:
– $\mathrm{Ag^+}\;\text{silver}$
– $\mathrm{Zn^{2+}}\;\text{zinc}$
– $\mathrm{Cd^{2+}}$ (implied although not on printed list)
• Molecular cation: $\mathrm{NH<em>4^+}\;\text{ammonium}$ • Di-nuclear: $\mathrm{Hg</em>2^{2+}}\;\text{mercury(I)}$

Type II (variable-charge) Cations — Roman numeral required

• $\mathrm{Fe^{3+}}$ iron(III); $\mathrm{Fe^{2+}}$ iron(II)
• $\mathrm{Cu^{2+}}$ copper(II); $\mathrm{Cu^{+}}$ copper(I)
• $\mathrm{Co^{3+}}$ cobalt(III); $\mathrm{Co^{2+}}$ cobalt(II)
• $\mathrm{Sn^{4+}}$ tin(IV); $\mathrm{Sn^{2+}}$ tin(II)
• $\mathrm{Pb^{4+}}$ lead(IV); $\mathrm{Pb^{2+}}$ lead(II)
• $\mathrm{Hg^{2+}}$ mercury(II) (note distinction from $\mathrm{Hg_2^{2+}}$)

Monatomic Anions (fixed charge)

• Group 17 (halogens, $-1$): $\mathrm{F^-}$ fluoride, $\mathrm{Cl^-}$ chloride, $\mathrm{Br^-}$ bromide, $\mathrm{I^-}$ iodide.
• Group 16 non-metals ($-2$): $\mathrm{O^{2-}}$ oxide, $\mathrm{S^{2-}}$ sulfide, $\mathrm{Se^{2-}}$ selenide.
• Group 15 non-metals ($-3$): $\mathrm{N^{3-}}$ nitride, $\mathrm{P^{3-}}$ phosphide, $\mathrm{As^{3-}}$ arsenide.
• Hydride $\mathrm{H^-}$ retains prefix “hydro” in acids.

Polyatomic Anions — Master List

• Oxo-ions
– $\mathrm{NO<em>3^-}$ nitrate ➜ $\mathrm{NO</em>2^-}$ nitrite
– $\mathrm{SO<em>4^{2-}}$ sulfate ➜ $\mathrm{SO</em>3^{2-}}$ sulfite
– $\mathrm{PO<em>4^{3-}}$ phosphate – $\mathrm{CO</em>3^{2-}}$ carbonate
– $\mathrm{CrO<em>4^{2-}}$ chromate ➜ $\mathrm{Cr</em>2O<em>7^{2-}}$ dichromate – $\mathrm{MnO</em>4^-}$ permanganate
– $\mathrm{BO<em>3^{3-}}$ borate – $\mathrm{C</em>2O<em>4^{2-}}$ oxalate – $\mathrm{S</em>2O<em>3^{2-}}$ thiosulfate • Hydrogen/bi- derivatives (charge increases by +1 per hydrogen): $\mathrm{HSO</em>4^-}$, $\mathrm{HPO<em>4^{2-}}$, $\mathrm{H</em>2PO<em>4^-}$, $\mathrm{HCO</em>3^-}$
• Halogen oxo-series — memorize one, transfer to Br, I:
$\mathrm{ClO^-}$ (hypo-), $\mathrm{ClO<em>2^-}$ (-ite), $\mathrm{ClO</em>3^-}$ (-ate), $\mathrm{ClO<em>4^-}$ (per-). • Other special anions – $\mathrm{OH^-}$ hydroxide – $\mathrm{CN^-}$ cyanide – $\mathrm{SCN^-}\text{ or }NCS^-$ thiocyanate – $\mathrm{C</em>2H<em>3O</em>2^-}$ or $\mathrm{CH<em>3COO^-}$ acetate – $\mathrm{O</em>2^{2-}}$ peroxide
– $\mathrm{NH_2^-}$ amide

Periodic-Table Charge Patterns & Logic

• Group 1: lose 1 e⁻ ⟶ $+1$ cations (alkali metals).
• Group 2: lose 2 e⁻ ⟶ $+2$ cations (alkaline earth).
• Group 13 metals (Al, Ga, In): lose 3 e⁻ ⟶ $+3$.
• Group 17: gain 1 e⁻ ⟶ $-1$.
• Group 16: gain 2 e⁻ ⟶ $-2$.
• Group 15: gain 3 e⁻ ⟶ $-3$.
• Naming rule: cations keep element name; anions take “-ide.”
• Variable-charge metals identified via Roman numeral in name (iron(II), tin(IV)…).

Polyatomic Anion Patterns & Mnemonics

• “-ate” vs “-ite”: same charge, but “-ite” has one fewer O.
Ex: sulfate $\rightarrow$ sulfite: $\mathrm{SO<em>4^{2-}\rightarrow SO</em>3^{2-}}$.
• Adding $\mathrm{H^+}$ prepends an H and raises the charge by +1:
$\mathrm{PO<em>4^{3-}\;\xrightarrow{+H^+}\;HPO</em>4^{2-}\;\xrightarrow{+H^+}\;H<em>2PO</em>4^-}$.
• Halogen ladder (prefix logic holds for BrO and IO series):
hypo- (-1 O) ⟶ “-ite” ⟶ “-ate” ⟶ per- ( +1 O).
Sequence retains charge throughout.
• Prefix etymology: “hypo” = under/too little (hypothermia); “hyper” ⟶ “per.”

Ionic Nomenclature Rules & Procedure

• Formula → Name

1. Identify cation; state its name.
2. Identify anion; state its name (monatomic = “-ide,” polyatomic = memorized).
3. For variable-charge metals, calculate charge from anion totals; insert Roman numeral.
   • Examples
   – $\mathrm{AlCl<em>3}$ ⟶ aluminum chloride. – $\mathrm{Na</em>2SO<em>4}$ ⟶ sodium sulfate. – $\mathrm{FeCl</em>3}$: three $\mathrm{Cl^-}$ = −3 ⇒ Fe = +3 ⟶ iron(III) chloride.
   – $\mathrm{Cr(NO<em>3)</em>3}$: three nitrates (−3) ⇒ Cr +3 ⟶ chromium(III) nitrate.
   • Name → Formula (electrostatic balancing)
4. Write cation symbol/charge.
5. Write anion symbol/charge.
6. Cross-multiply charges to smallest common multiple; polyatomics get parentheses.
   Ex: calcium phosphide: $\mathrm{Ca^{2+}}$ & $\mathrm{P^{3-}}$ ⟶ $\mathrm{Ca<em>3P</em>2}$.
   Ex: aluminum nitrite: $\mathrm{Al^{3+}}$ & $\mathrm{NO<em>2^-}$ ⟶ $\mathrm{Al(NO</em>2)_3}$.

Practice List (from worksheet)

• You must be able to complete conversions such as:
PbCl2, Na2C2O4, Cu(NO3)2, MgO2, Ca(C2H3O2)2, etc.

Binary Molecular (Covalent) Nomenclature

• Applies exclusively to two non-metals.
• First element: retains name; prefix only if subscript > 1.
• Second element: “-ide” ending & ALWAYS a prefix.
• Greek prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.
• Examples (name → formula)
– carbon dioxide ⟶ $\mathrm{CO<em>2}$ – dinitrogen monoxide ⟶ $\mathrm{N</em>2O}$
– phosphorus pentafluoride ⟶ $\mathrm{PF<em>5}$ • Examples (formula → name) – $\mathrm{N</em>2O<em>4}$ ⟶ dinitrogen tetroxide. – $\mathrm{SO</em>3}$ ⟶ sulfur trioxide.

Acid Naming Rules

Binary Acids (HX)

• Format: “hydro” + root of non-metal + “ic” + acid.
– $\mathrm{HCl}$ ⟶ hydrochloric acid
– $\mathrm{HF}$ ⟶ hydrofluoric acid

Oxyacids (contain O)

• Reference point: anions ending in “-ate” produce “-ic” acids.
$\mathrm{SO<em>4^{2-}}$ (sulfate) ⟶ $\mathrm{H</em>2SO<em>4}$ sulfuric acid. • One more O than “-ic” ⇒ prefix “per-” + “-ic.” $\mathrm{H</em>2SO<em>5}$ persulfuric acid. • One fewer O than “-ic” ⇒ change “-ic” to “-ous.” $\mathrm{H</em>2SO<em>3}$ sulfurous acid. • Two fewer O than “-ic” (one fewer than “-ous”) ⇒ prefix “hypo-” + “-ous.” $\mathrm{H</em>2SO_2}$ hyposulfurous acid.
• Same ladder applies to nitrate, chlorate, bromate, phosphate, carbonate, etc.

Quick Acid Naming Algorithm

1. Identify anion name.
2. Replace
   – “-ate” → “-ic.”
   – “-ite” → “-ous.”
3. Add prefixes per- or hypo- as oxygen count demands.
4. Append “acid.”

Practice Prompts (from worksheet)

• Binary: HF, HCl, H2S, HBr, HI.
• Oxyacids: sequence for carbonates, chlorates, phosphates given (must name all).
• Reverse problems: supply formulas for perbromic, nitrous, hypobromous, chromic, chromous, etc.

Ethical & Practical Implications

• Precise nomenclature prevents laboratory accidents (e.g., mixing nitrate vs nitrite).
• Pharmaceutical formulations rely on correct ion forms (e.g., $\mathrm{Fe^{2+}}$ vs $\mathrm{Fe^{3+}}$ bioavailability).
• Environmental chemistry distinguishes sulfate vs sulfite in atmospheric studies of acid rain.

Real-World Connections & Mnemonics

• Medical: “hypo” in hypoglycemia mirrors “hypo-” in hypochlorite (shortage/under).
• Industrial bleaching: $\mathrm{ClO^-}$ hypochlorite is the active ion in household bleach.
• Car batteries: lead(IV) oxide vs lead(II) sulfate highlight variable oxidation states.

Summary of Key Equations & Relationships

• Charge balance rule: $\sum q<em>+ + \sum q</em>- = 0$ in any neutral compound.
• Hydrogen addition: $\mathrm{X^{n-}} + \mathrm{H^+} \rightarrow H\mathrm{X}^{(n-1)-}$ (iterative logic for bi-/di- hydrogen ions).
• Oxygen ladder for halogens: $\mathrm{XO^-}\;(\text{hypo}) \rightarrow XO<em>2^- \rightarrow XO</em>3^- \rightarrow XO_4^-\;(\text{per})$ (charge constant).

Memorization Tip: Create flashcards grouping ions by family (e.g., all “ites/ates” of nitrogen, all variable-charge metals) and practice daily until recall is instantaneous.