Chemistry Notes: Ionic & Covalent Compounds, Formulas, and Moles

Ionic Compounds & Polyatomic Ions

• Meeting context: Chemistry class session led by an instructor with Luke participating and asking questions; covered ionic and covalent compounds, formulas, naming conventions, and calculations involving moles and atomic mass; included Q&A on course logistics, grading, and homework resources.
• Calcium phosphate formula: calcium (Ca^{2+}) and phosphate (PO4^{3-}) ions combine so charges add to zero. Balancing example: 3 Ca^{2+} (total +6) with 2 PO4^{3-} (total -6) yields a neutral compound.
• Use parentheses for polyatomic ions in formulas (e.g., $\mathrm{Ca<em>3(PO</em>4)_2}$).
• Polyatomic ions: clusters of atoms with a net charge, can be positive or negative.
• Only need to memorize common polyatomic ions for exams/homework.
• Ionic compounds: always balance charges to zero; cation written first, anion second.

Covalent Bonding & Molecules

• Covalent bonds form when attractive forces (nucleus–electron) outweigh repulsive forces (nucleus–nucleus, electron–electron).
• Covalent bonds: electrons are shared between nonmetals.
• Molecules: two or more atoms bonded together; can be elements (e.g., $\mathrm{O<em>2}$) or compounds (e.g., $\mathrm{H</em>2O}$).
• Diatomic elements to memorize: $\mathrm{H<em>2}, \mathrm{O</em>2}, \mathrm{N<em>2}, \mathrm{Cl</em>2}, \mathrm{Br<em>2}, \mathrm{I</em>2}, \mathrm{F_2}$.
• Diatomic elements are always found as pairs in nature; mnemonic provided in class slides (reference to class resource).

Formulas: Empirical, Molecular, Structural

• Empirical formula: simplest whole-number ratio of atoms in a substance (e.g., $\mathrm{CH_2O}$ for glucose).
• Molecular formula: actual number of each atom in a molecule (e.g., $\mathrm{C<em>6H</em>{12}O_6}$ for glucose).
• Structural formula: shows how atoms are connected (e.g., $\mathrm{H-O-H}$ for water).
• Condensed structural formulas: shorthand for connectivity, used in organic chemistry.
• Some compounds (like $\mathrm{H_2O}$) have the same empirical and molecular formula.

Naming Compounds

• Ionic compounds: no prefixes; use empirical formula; cation written first, anion second.
• Covalent compounds: use prefixes (mono-, di-, tri-, tetra-, penta-, hexa-, etc.) for the number of atoms.

Page 2 – Acids, Special Cases & Calculations

• Binary compounds: only two elements; water and ammonia are exceptions with common names (e.g., water is common name for $\mathrm{H<em>2O}$; ammonia for $\mathrm{NH</em>3}$).
• Acids: hydrogen + nonmetal; named differently in water (e.g., hydrochloric acid for HCl) and in other contexts.
• Acids are covalent but produce ions in water;
  • Naming depends on whether the acid is in water (aqueous) or not.
• Strong acids discussed: $\mathrm{HCl}, \mathrm{HBr}, \mathrm{HI}, \mathrm{HF}$.
• H^+ ion is just a proton (no electrons or neutrons): $\mathrm{H^+}$.
• Calculations: Moles, Atomic Mass, Conversions
  • Mole: standard counting unit; $1\ \text{mol} = N_A = 6.022 \times 10^{23}$ particles (Avogadro’s number).
  • Formula: $N_A = 6.022 \times 10^{23}$.
  • Molar mass: mass of 1 mole of a substance, in $\text{g/mol}$.
  • Use conversion factors to go between atoms, moles, and grams.
  • Always specify units (atoms, molecules, formula units) in calculations.
  • Practice problems: converting atoms to moles, grams to moles, etc.

Page 2 – Course Logistics & Grading

• Canvas gradebook doesn’t show running total; use a grade calculator for estimates.
• Attendance and participation points tracked separately.
• Homework and quizzes: practice with provided resources, prefixes, and naming conventions.

Page 2 – Study Tips & Practice

• Use provided lists for polyatomic ions and acids as references.
• Prefixes for covalent naming: memorize up to six (mono-, di-, tri-, tetra-, penta-, hexa-).
• Practice converting between different types of formulas and units.
• Instructor encourages making mistakes now, not on tests.

Page 2 – Sidebar & Miscellaneous

• Organic chemistry and polymers briefly mentioned as extensions.
• Polymers: repeating units (monomers) form large molecules (e.g., polyethylene).
• Instructor available for office hours for further questions or test review.

Page 3 – Key Takeaways

• Know how to write and name ionic and covalent compounds.
• Understand differences between empirical, molecular, and structural formulas.
• Be comfortable with mole calculations and unit conversions.
• Use provided resources and ask questions as needed for clarification.

Quick reference formulas and concepts

• Calcium phosphate balance:
  • Cations: $\mathrm{Ca^{2+}}$
  • Anions: $\mathrm{PO_4^{3-}}$
  • Net neutral example: $3\mathrm{Ca^{2+}} + 2\mathrm{PO<em>4^{3-}} \rightarrow \mathrm{Ca</em>3(PO<em>4)</em>2}$
• Polyatomic ions: memorize common ions (e.g., $\mathrm{SO<em>4^{2-}}, \mathrm{NO</em>3^{-}}, \mathrm{OH^{-}}$, etc.); not all listed here, but memorize as needed for coursework.
• Empirical vs. molecular: empirical shows simplest ratio; molecular shows actual counts (e.g., glucose: empirical $\mathrm{CH<em>2O}$; molecular $\mathrm{C</em>6H<em>{12}O</em>6}$).
• Structural: connectivities (e.g., water: $\mathrm{H-O-H}$).
• Prefixes in covalent naming: $\text{mono-}, \text{di-}, \text{tri-}, \text{tetra-}, \text{penta-}, \text{hexa-}$
• Diatomic elements (seven): $\mathrm{H<em>2}, \mathrm{O</em>2}, \mathrm{N<em>2}, \mathrm{Cl</em>2}, \mathrm{Br<em>2}, \mathrm{I</em>2}, \mathrm{F_2}$
• Avogadro’s number: $N<em>A = 6.022 \times 10^{23}$; 1 mole contains $N</em>A$ particles.
• Molar mass unit: $M \ ( ext{g/mol})$; conversion between grams and moles: $n = \frac{m}{M}$; atoms to moles: $n = \frac{N}{N_A}$; moles to grams: $m = nM$.