Introduction to Ionic and Covalent Compounds, Naming, Stoichiometry, and Thermochemistry

Chemical Formulas and Compound Types

• Chemical Formulas: Used to represent compounds. They consist of:

  • Symbols: Represent the elements involved.
  • Subscripts: Determine the number of atoms of a specific element.
  • Parentheses: Used to contain groups of atoms (polyatomic ions) that act as a single unit.

• Formula Examples and Atomic Counts:

  • $H_2S$: $2$ atoms of $H$, $1$ atom of $S$.
  • $Cl_2$: $2$ atoms of $Cl$.
  • $NaNO_3$: $1$ atom of $Na$, $1$ atom of $N$, $3$ atoms of $O$.
  • $Pb(NO_3)_2$: $1$ atom of $Pb$, $2$ atoms of $N$, $6$ atoms of $O$.
  • $(NH_4)_2SO_4$: $2$ atoms of $N$, $8$ atoms of $H$, $1$ atom of $S$, $4$ atoms of $O$.

• Classification of Compounds:

  • Covalent Compounds: Contain covalent bonds. They typically involve nonmetals bonded to other nonmetals.
    • Examples: Water ($H_2O$), Carbon dioxide ($CO_2$), Glucose ($C_6H_12O_6$).
  • Ionic Compounds: Contain ionic bonds. They involve nonmetals bonded to metals.
    • Examples: Table salt ($NaCl$), Baking soda ($NaHCO_3$), Milk of magnesia ($Mg(OH)_2$), Rubies and Sapphires ($Al_2O_3$ with trace metals).

• Significance of Compound Type: The type of compound determines physical properties critical for medicinal use, such as:

  • Medicinal purpose.
  • Solubility (many drugs are formulated as salts to increase water solubility).
  • Storage stability and formulation methods.

Ion Formation and Charge

• Cations (Positive Ions):

  • Formed by metals losing enough electrons to achieve the same electron count as the noble gas in the row above them.
  • Group 1A (1): Form $1+$ ions.
  • Group 2A (2): Form $2+$ ions.
  • Group 3A (3): Form $3+$ ions.
  • Sodium Ion ($Na^+$): Sodium atom ($Na$ with $11\text{ p}^+$ and $11\text{ e}^-$) loses $1$ valence electron to become sodium ion ($11\text{ p}^+$ and $10\text{ e}^-$).
  • Magnesium Ion ($Mg^{2+}$): Magnesium atom ($Mg$ with $12\text{ p}^+$ and $12\text{ e}^-$) loses $2$ valence electrons to become magnesium ion ($12\text{ p}^+$ and $10\text{ e}^-$).

• Anions (Negative Ions):

  • Formed by nonmetals gaining enough electrons to match the noble gas in the same row of the periodic table.
  • Group 7A (17): Form $1^-$ ions.
  • Group 6A (16): Form $2^-$ ions.
  • Group 5A (15): Form $3^-$ ions.
  • Chloride Ion ($Cl^-$): Chlorine atom ($Cl$ with $17\text{ p}^+$ and $17\text{ e}^-$) gains $1$ valence electron to become chloride ion ($17\text{ p}^+$ and $18\text{ e}^-$).

• Common Ionic Charges Table Summary:

  • $Li^+$, $Na^+$, $K^+$: $+1$
  • $Mg^{2+}$, $Ca^{2+}$, $Zn^{2+}$: $+2$
  • $Al^{3+}$: $+3$
  • $N^{3-}$, $P^{3-}$: $-3$
  • $O^{2-}$, $S^{2-}$: $-2$
  • $F^-$, $Cl^-$, $Br^-$, $I^-$: $-1$
  • Variable Charge Metals: $Cr^{2+}/Cr^{3+}$, $Mn^{2+}/Mn^{3+}$, $Fe^{2+}/Fe^{3+}$, $Co^{2+}/Co^{3+}$, $Ni^{2+}/Ni^{3+}$, $Cu^+/Cu^{2+}$, $Au^+/Au^{3+}$, $Hg_2^{2+}/Hg^{2+}$, $Sn^{2+}/Sn^{4+}$, $Pb^{2+}/Pb^{4+}$.

Ionic Compounds and Nomenclature

• Properties of Ionic Compounds:

  • Formed by a cation and an anion attracted by opposite charges (ionic bond).
  • Strong bonds resulting in high melting and boiling points (very stable).
  • Typically solids at room temperature.
  • Salts: Common term for ionic compounds. Examples include medicinal salts like Phenylephrine $HCl$ (nasal decongestant) and Dextromethorphan $HBr$ (cough suppressant).

• Formula Writing Rules:

  • Cation is always written first.
  • Uses subscripts to represent the lowest whole-number ratio of ions.
  • Subscripts must balance the charges to reach a net charge of zero.
  • Example $Na_2S$: Sulfur carries a $2^-$ charge; two Sodium ions (each $1^+$) are required to balance it: $2(1^+) + 1(2^-) = 0$.

• Polyatomic Ions: Groups of covalently bonded atoms carrying a collective charge.

  • Ammonium: $NH_4^+$
  • Hydroxide: $OH^-$
  • Nitrate: $NO_3^-$
  • Nitrite: $NO_2^-$
  • Carbonate: $CO_3^{2-}$
  • Phosphate: $PO_4^{3-}$
  • Hydrogen Carbonate (Bicarbonate): $HCO_3^-$
  • Rule: Use parentheses when more than one polyatomic ion is needed in a formula (e.g., $(NH_4)_2S$).

• Naming Single-Atom Ions:

  • Cations: Retain the element name (e.g., Sodium ion).
  • Anions: End in "-ide" (e.g., Chloride).

• Naming Transition Metals: Must use a Roman numeral in parentheses to indicate the specific charge (e.g., $FeCl_3$ is Iron (III) chloride).

• Polyatomic Naming Strategies:

  • The Big 7: Polyatomic anions with $3$ oxygen atoms ending in "-ate" ($NO_3^-$ Nitrate, $ClO_3^-$ Chlorate, etc.).
  • The Little 7: Polyatomic anions ending in "-ite" like Phosphite ($PO_3^{3-}$) and Sulfite ($SO_3^{2-}$).

• Hydrates: Compounds containing trapped water molecules in the crystal structure.

  • Named as the ionic compound + [prefix]-hydrate.
  • Prefixes: $1$ = mono, $2$ = di, $3$ = tri, $4$ = tetra, $5$ = penta, $6$ = hexa, $7$ = hepta, $8$ = octa, $9$ = nona, $10$ = deca.
  • Example: $CuSO_4 \cdot 5H_2O$ is Copper (II) sulfate pentahydrate.

Covalent Compounds and Acids Nomenclature

• Binary Covalent Compounds:
  • First element retains its original name.
  • Second element ends in "-ide".
  • Use prefixes to indicate the number of atoms (mono-, di-, tri-, etc.).
  • Exception: