Chemical Formulas, Ion Charges, and Types of Compounds

Chemical Formulas and Ion Charges

Understanding Ion Charges

When forming chemical compounds, it is essential that the total charges of all ions in the compound balance to zero. This ensures the compound is electrically neutral.

• Ions: Atoms that gain or lose electrons to achieve a stable electron configuration.

  • Cation: A positively charged ion formed by losing electrons.

  • Anion: A negatively charged ion formed by gaining electrons.

• Example:

  • Lead (Pb) forms a +2 charge cation.

  • Chlorine (Cl) forms a –1 charge anion.

  • Formula: PbCl2 (charges balance: +2 and two –1s)

Determining Ionic Formulas

To write the correct formula for an ionic compound, follow these steps:

1. Identify each element and its group number on the periodic table.

2. Determine the charge based on the group:

   • Group 1 → +1

   • Group 2 → +2

   • Group 15 → –3

   • Group 16 → –2

   • Group 17 → –1

3. Combine ions in ratios that balance the total positive and negative charges to zero, adjusting subscripts as needed.

4. Example: Calcium (Ca2+) and Nitrogen (N3–): Cross-multiply charges to get Ca3N2.

Ionic vs. Covalent Compounds

Ionic Compounds

Ionic compounds are formed by the transfer of electrons from metals to nonmetals, resulting in the formation of cations and anions. These ions arrange in repeating patterns to create crystalline structures.

• Formed between metals and nonmetals

• Electrons are transferred

• Create crystalline structures (lattices of alternating positive and negative ions)

• Example: NaCl (sodium chloride)

Covalent Compounds

Covalent compounds are formed when nonmetals share electrons to achieve stable electron configurations. These compounds exist as discrete molecules rather than extended lattices.

• Formed between nonmetals

• Electrons are shared

• Form molecules, not crystals

• Examples: H2O, CO2, NCl3

Example: Magnesium Oxide (MgO)

• Magnesium (Group 2) loses 2 electrons → Mg2+

• Oxygen (Group 16) gains 2 electrons → O2–

• Combine to form MgO (charges balance: +2 and –2)

Examples of Covalent Compounds

Example: Nitrogen Trichloride (NCl3)

• Both nitrogen and chlorine are nonmetals → covalent compound

• Nitrogen: 5 valence electrons (Group 15)

• Each Chlorine: 7 valence electrons (Group 17)

• Total electrons = 5 + (3 × 7) = 26

• Central atom: Nitrogen

• Each Cl forms a single bond with N

• Remaining electrons form lone pairs to complete octets

• Final structure: N single-bonded to 3 Cl atoms

Covalent Bond Behavior

• Molecules exist independently and move freely (especially in liquids and gases).

• Enormous numbers of molecules can exist in small amounts of material (around $10^{22}$ molecules per drop — Avogadro's number).

Example: Carbonate Ion (CO32–)

• Found in: blood, bones, rocks, rivers, and shells (like limestone and clams)

• Carbonate = CO32–

• Carbon (4 valence electrons), 3 Oxygens (6 each), and extra 2 electrons from charge

• Central atom: Carbon

• Arrange 3 oxygens around it

• Use bonds (6 electrons), then distribute remaining as lone pairs

• One oxygen forms a double bond with carbon

• Final structure satisfies octets for all atoms

• Add –2 charge to the ion (CO32–)

Study Advice from Instructor

• Do not rely on memorization — chemistry requires understanding and skills.

• Practice building compounds and balancing charges.

• Use provided worksheets for skill-building.

• Review why certain formulas form the way they do — not just what they are.

• Many students succeeded by learning the process, not just the answers.