Polyatomic Ions and Ionic Compounds - Study Notes

Polyatomic Ions: Overview

• Ions can be composed of groups of atoms covalently bonded together; they have an overall electric charge. When a group of atoms behaves as a single unit with a charge, it is called a polyatomic ion (poly = more than one atom).

• Polyatomic ions can be anions (negative charge) or cations (positive charge).

• These ions form a wide variety of ionic compounds when combined with other ions.

Polyatomic Anions

• Nitrogen-centered:

  • Nitrite: $NO_2^-$

  • Nitrate: $NO_3^-$

  • Both have an overall charge of $-1$ and differ in the number of oxygen atoms (nitrite has 2 O, nitrate has 3 O).

• Sulfur-centered:

  • Sulfite: $SO_3^{2-}$

  • Sulfate: $SO_4^{2-}$

  • Both have an overall charge of $-2$; differ by one oxygen atom (3 vs 4).

  • Example: Magnesium sulfate: $MgSO_4$, used as a laxative.

• Phosphorus-centered:

  • Phosphate: $PO_4^{3-}$ (one phosphorus, four oxygens; charge $-3$)

  • Hydrogen phosphate: $HPO_4^{2-}$ (one hydrogen, one phosphorus, four oxygens; charge $-2$)

  • Dihydrogen phosphate: $H<em>2PO</em>4^{-}$ (two hydrogens, one phosphorus, four oxygens; charge $-1$)

  • These phosphate-related ions are important for buffering and biological processes; phosphate ions form a wide variety of ionic compounds (e.g., sodium phosphate as a fertilizer).

• Carbon-centered:

  • Carbonate: $CO_3^{2-}$

  • Hydrogen carbonate (bicarbonate): $HCO_3^{-}$; hydrogen carbonate is also called bicarbonate.

  • Sodium hydrogen carbonate: $NaHCO_3$ (baking soda; antacid).

  • Carbonate ions form many compounds; carbonates are found in chalk and eggshells, and are used in antacids (e.g., TUMS).

• Hydrogen carbonate continues to emphasize the bicarbonate role as a buffer in biological systems and food chemistry.

• Hydroxide:

  • Hydroxide: $OH^-$

  • One oxygen and one hydrogen; charge $-1$; forms many ionic hydroxide compounds (e.g., magnesium hydroxide).

  • Magnesium hydroxide: $Mg(OH)_2$ (active ingredient in milk of magnesia; antacid).

• Acetate (brief note):

  • Acetate ion is a polyatomic anion; typically $CH<em>3COO^-$ or $CH</em>3COO^-$ depending on context. (Mentioned as important, to be covered in Chapter 4 with organic chemistry intro.)

Polyatomic Cations

• Hydronium ion:

  • Hydronium: $H_3O^+$

  • Formed when water accepts a proton; a polyatomic cation.

  • Important in acid-base chemistry and pH discussions (Unit 3).

• Ammonium ion:

  • Ammonium: $NH_4^+$

  • Ammonia $NH_3$ accepts a proton to form ammonium; Ammonium salts (e.g., ammonium phosphate, ammonium sulfate, ammonium nitrate) are widely used as fertilizers.

• Both hydronium and ammonium are polyatomic cations with one positive charge.

Ionic Compounds Containing Polyatomic Ions

• Polyatomic ions act as a unit in ionic compounds (they don’t break apart in the formula).

• Example: Sodium sulfate

  • Cation: $Na^+$; Anion: $SO_4^{2-}$

  • Neutrality requires two $Na^+$ for one $SO<em>4^{2-}$: formula $Na</em>2SO_4$; name: sodium sulfate.

• Magnesium hydroxide

  • One $Mg^{2+}$ and two $OH^-$; to show two hydroxide ions, brackets are used: $Mg(OH)_2$.

• Ammonium carbonate

  • Two ammonium ions for one carbonate: $(NH<em>4)</em>2CO_3$; name: ammonium carbonate.

• Transition metal cations require charge indication in names (Roman numerals) and appropriate subscripts in formulas.

  • Ferrous sulfate: $FeSO_4$ (Fe^{2+}); name reflects +2 charge (ferrous).

  • Ferric sulfate: $Fe<em>2(SO</em>4)<em>3$ (Fe^{3+}); criss-cross method used to balance charges; because sulfate is $SO</em>4^{2-}$, two Fe^{3+} ions balance three sulfate ions.

  • Copper carbonate or cupric carbonate: $CuCO_3$ (Cu^{2+}); ratio 1:1 with carbonate.

• Criss-cross method and bracket usage:

  • When there are multiple polyatomic ions, use brackets to indicate the number of polyatomic ions, e.g., $Fe<em>2(SO</em>4)<em>3$, ${(NH</em>4)<em>2CO</em>3}$, $Mg(OH)_2$.

  • If there is only a single polyatomic ion, brackets are not necessary (e.g., $FeSO_4$).

• Quick naming rules recap:

  • Cation is written first, then anion.

  • For transition metals, include the Roman numeral for charge if needed.

  • Use brackets around polyatomic ions when there are multiple of that ion in the formula.

Uses of Ionic Compounds in Everyday Life and Medicine

• Sodium iodide: $NaI$; source of iodide for thyroid.

• Sodium nitrite: $NaNO_2$; meat preservative.

• Sodium nitrate: $NaNO_3$; dairy/food preservative.

• Sodium hydrogen carbonate (baking soda): $NaHCO_3$; antacid.

• Calcium carbonate: $CaCO_3$; chalk; eggshell; used as an antacid (e.g., TUMS).

• Magnesium hydroxide: $Mg(OH)_2$; antacid (milk of magnesia).

• Barium sulfate: $BaSO_4$; extremely insoluble; used as barium milk in X-ray imaging of the GI tract.

• Calcium sulfate: $CaSO_4$; insoluble; used in plaster casts (plaster of Paris).

• Magnesium sulfate: $MgSO_4$; laxative.

• Silver nitrate: $AgNO_3$; tests for halide ions; antiseptic for newborn eye infections.

• Ammonium carbonate: $(NH<em>4)</em>2CO_3$; used as smelling salt due to strong odor.

• Sodium hydroxide: $NaOH$; caustic; used in drain cleaners.

• Lithium carbonate: $Li<em>2CO</em>3$; used in treatment of manic depression.

Ionic Compounds and Conductivity: Electrolytes

• Electrolytes are compounds that dissolve in water to form ions and conduct electricity.

• Demonstration concepts (three beakers):

  • Distilled water: pure water; non-conductor (no ions).

  • Solid ionic compound (e.g., table salt) in solid form: ions fixed in lattice; non-conductor.

  • A solution of an ionic compound in water: cations move to the negative electrode, anions move to the positive electrode; conducts electricity.

• Strong electrolytes vs weak electrolytes vs non-electrolytes:

  • Strong electrolytes: large number of ions in solution; good conductors (e.g., NaCl, HCl).

  • Weak electrolytes: few ions in solution; partial conductivity (e.g., vinegar, acetic acid, CH_3COOH in water).

  • Non-electrolytes: little to no ions in solution; poor or no conductivity (e.g., sucrose, C12H22O_11; rubbing alcohol like propanol).

• Biological relevance of electrolytes:

  • Sodium (Na^+) and potassium (K^+) ions are essential for nerve impulse transmission.

  • Calcium (Ca^{2+}) ions are important for blood clotting and muscle contraction.

• Demonstration examples (conductivity observations):

  • Sodium chloride solution (NaCl in water): strong electrolyte; conducts electricity well; isotonic at 0.95 ext{A0% NaCl} solution.

  • Sucrose solution: non-electrolyte; no conduction.

  • Vinegar (5% acetic acid, CH_3COOH in water): weak electrolyte; some conduction due to some ions, but not as strong as a strong electrolyte.

  • Hydrochloric acid (HCl) solution: very strong electrolyte; conducts electricity strongly.

  • Propanol (isopropanol): non-electrolyte; does not conduct electricity.

• Practical implications: electrolytes are critical in health care, especially for hydration and bodily function; dehydration treatment often involves electrolyte solutions.

Study Tips for Polyatomic Ions

• Flashcard strategy (six cards):

  • Card 1: Nitrite $NO<em>2^-$ and Nitrate $NO</em>3^-$ (centered around nitrogen).

  • Card 2: Sulfite $SO<em>3^{2-}$ and Sulfate $SO</em>4^{2-}$ (centered around sulfur).

  • Card 3: Phosphate $PO<em>4^{3-}$, Hydrogen phosphate $HPO</em>4^{2-}$, Dihydrogen phosphate $H<em>2PO</em>4^{-}$ (centered around phosphorus).

  • Card 4: Carbonate $CO<em>3^{2-}$, Hydrogen carbonate $HCO</em>3^{-}$ (bicarbonate).

  • Card 5: Hydroxide $OH^-$.

  • Card 6: Hydronium $H<em>3O^+$ and Ammonium $NH</em>4^+$.

• Study approach: review these cards two or three times daily to memorize names, formulas, and charges.

Quick Reference: Common Formulas and Names (selected)

• Nitrite: $NO<em>2^-$; Nitrate: $NO</em>3^-$

• Sulfite: $SO<em>3^{2-}$; Sulfate: $SO</em>4^{2-}$

• Phosphate: $PO<em>4^{3-}$; Hydrogen phosphate: $HPO</em>4^{2-}$; Dihydrogen phosphate: $H<em>2PO</em>4^{-}$

• Carbonate: $CO<em>3^{2-}$; Hydrogen carbonate: $HCO</em>3^{-}$

• Hydroxide: $OH^-$

• Hydronium: $H<em>3O^+$; Ammonium: $NH</em>4^+$

• Sodium sulfate: $Na<em>2SO</em>4$; Magnesium hydroxide: $Mg(OH)<em>2$; Ammonium carbonate: $(NH</em>4)<em>2CO</em>3$

• Ferrous sulfate: $FeSO<em>4$; Ferric sulfate: $Fe</em>2(SO<em>4)</em>3$; Copper carbonate: $CuCO_3$

• Sodium iodide: $NaI$; Sodium nitrite: $NaNO<em>2$; Sodium nitrate: $NaNO</em>3$

• Sodium bicarbonate: $NaHCO<em>3$; Calcium carbonate: $CaCO</em>3$; Magnesium sulfate: $MgSO<em>4$; Barium sulfate: $BaSO</em>4$; Calcium sulfate: $CaSO<em>4$; Magnesium hydroxide: $Mg(OH)</em>2$; Silver nitrate: $AgNO<em>3$; Ammonium carbonate: $(NH</em>4)<em>2CO</em>3$; Sodium hydroxide: $NaOH$; Lithium carbonate: $Li<em>2CO</em>3$

Notes on Formulas and Naming

• When more than one polyatomic ion is present in an ionic compound, use brackets to indicate the quantity, e.g., $Mg(OH)<em>2$, ${(NH</em>4)<em>2CO</em>3}$, $Fe<em>2(SO</em>4)_3$.

• If there is only a single polyatomic ion, brackets are not necessary, e.g., $FeSO_4$ (ferrous sulfate).

• The cation always precedes the anion in the formula and name; transition metals use Roman numerals to indicate charge when necessary.

• The term criss-cross method helps balance charges when writing formulas for ionic compounds (e.g., $Fe^{3+}$ with $SO<em>4^{2-}$ yields $Fe</em>2(SO<em>4)</em>3$).

Reminder: Key Concepts to Remember

• Polyatomic ions behave as units in compounds.

• Different oxyanion forms of a given element have different charges and O counts (nitrogen, sulfur, phosphorus, carbon groups).

• Brackets indicate multiple polyatomic ions in a formula; single polyatomic ions do not require brackets.

• Strong electrolytes vs weak electrolytes vs non-electrolytes determine electrical conductivity in solution.

• Biological relevance of ions (Na^+, K^+, Ca^{2+}, H_3O^+) and practical medical uses of ionic compounds.