Chapter 2: Ionic and Molecular Compounds

Classification of Chemical Compounds

• General Classification: Compounds in chemistry are broadly classified into two main types: ionic compounds and molecular (covalent) compounds.

• Ionic Compounds: These are formed by the combination of metals and nonmetals.

  • In these compounds, the metal acts as a cation (a positively charged ion).

  • The nonmetal acts as an anion (a negatively charged ion).

• Molecular (Covalent) Compounds: These are formed between two nonmetals or between a nonmetal and a metalloid.

  • In these compounds, both elements are typically ones that can form anions (negative charges), though they share electrons rather than transferring them.

The Periodic Table and Ion Formation

• Metals: Highlighted in green on the periodic table. These elements are known to form cations (positive charge).

• Nonmetals: Highlighted in blue on the periodic table, including Hydrogen. These elements are known to form anions (negative charge).

• Metalloids: Highlighted in grey on the periodic table.

• Representative Elements: These are elements found in Groups 1 through 8. Their charge states are often predictable based on their group number.

• Cation Charges (Metals):

  • Group 1: Elements in this group form a $+1$ cation.

  • Group 2: Elements in this group form a $+2$ cation.

  • Group 3: Elements in this group can form a $+3$ cation.

• Anion Charges (Nonmetals):

  • The goal for nonmetals in Group 5 through Group 7 is to achieve a total of 8 electrons (stability).

  • Group 5: Requires 3 more electrons to reach 8. They form a $-3$ anion.

  • Group 6: Requires 2 more electrons to reach 8. They form a $-2$ anion.

  • Group 7: Requires 1 more electron to reach 8. They form a $-1$ anion.

• Exceptions and Omissions:

  • Group 8 (Noble Gases): These are stable and are generally left out of compound formation discussions at this level.

  • Group 4: This group is omitted from the general discussion of ionic and molecular compounds because its first member, Carbon, involves a specialized field known as organic chemistry.

• Transition Metals: These are the elements located between Group 2 and Group 3.

  • Unlike representative metals, transition metals can form more than one kind of cation charge.

  • Example: Chromium ($Cr$) can form a $+2$ and a $-3$ cation (as stated in the transcript).

Foundational Criteria for Ionic Formulas

Every ionic compound formula must meet two specific criteria to be considered correct:

1. Empirical Form: The molecular formula of an ionic compound is usually its empirical formula, which is the simplest form of the ratio that cannot be reduced further.

   • Illustration: If a formula is $X_2Y_4$, it is not empirical. It must be reduced by a factor of 2 to become $XY_2$.

2. Zero Net Charge: The sum of the charges of the cations and anions in each formula unit must equal exactly zero. If they do not add up to zero, the formula is incorrect.

Writing and Verifying Ionic Formulas

• The Cross-Over Rule: When a cation and anion combine, the coefficient of the charge of the cation becomes the subscript of the anion, and the coefficient of the charge of the anion becomes the subscript of the cation.

• Example: Sodium Chloride ($NaCl$):

  1. Sodium ($Na$) is in Group 1 ($+1$ charge).

  2. Chlorine ($Cl$) is in Group 7 ($-1$ charge).

  3. Formula development: $1$ $Na$ to $1$ $Cl$ ratio. In chemistry, if the subscript is $1$, it is not written.

  4. Check Criterion 1: Ratio is $1:1$, which is empirical.

  5. Check Criterion 2: $(1 \times (+1)) + (1 \times (-1)) = 0$. The formula is correct.

• Example: Aluminum Oxide ($Al_2O_3$):

  1. Aluminum ($Al$) is in Group 3 ($+3$ charge).

  2. Oxygen ($O$) is in Group 6 ($-2$ charge).

  3. Formula: $Al_2O_3$.

  4. Check Criterion 1: $2:3$ ratio cannot be reduced; it is empirical.

  5. Check Criterion 2: $(2 \times (+3)) + (3 \times (-2)) = +6 - 6 = 0$. The formula is correct.

• Example: Calcium Bromide ($CaBr_2$):

  1. Calcium ($Ca$) is in Group 2 ($+2$ charge).

  2. Bromine ($Br$) is in Group 7 ($-1$ charge).

  3. Formula: $CaBr_2$.

  4. Check Criterion 1: $1:2$ ratio is empirical.

  5. Check Criterion 2: $(1 \times (+2)) + (2 \times (-1)) = +2 - 2 = 0$. The formula is correct.

Polyatomic Ions and Ionic Formulas

• Definition: Polyatomic ions consist of multiple atoms acting as a single charged unit. These cannot be found directly on the periodic table; a reference table is usually required.

• Monoatomic Ions: Simple ions consisting of a single element from the periodic table.

• Use of Parentheses: Parentheses are required in a formula whenever a polyatomic ion has a subscript greater than $1$. This indicates that the subscript applies to the entire polyatomic unit, not just the last atom listed.

  • Example: Magnesium Hydroxide ($Mg(OH)_2$): Magnesium ($Mg^{2+}$) combines with Hydroxide ($OH^-$). The $2$ from Magnesium applies to the entire $OH$ unit. Writing it as $MgOH_2$ would be incorrect as the $2$ would only apply to the Hydrogen.

  • Example: Sodium Carbonate ($Na_2CO_3$): Sodium ($Na^+$) and Carbonate ($CO_3^{2-}$). The formula is $Na_2CO_3$. Since Carbonate has a subscript of $1$, parentheses are not needed.

  • Calculation for $Na_2CO_3$: $(2 \times (+1)) + (1 \times (-2)) = 0$.

Nomenclature of Ionic Compounds

• General Rule: The metal (cation) is named first, followed by the nonmetal (anion). The nonmetal name is modified to end with the suffix -IDE.

  • $BaCl_2$: Barium chloride.

  • $K_2O$: Potassium oxide.

  • $Mg(OH)_2$: Magnesium hydroxide (no "ide" added to the polyatomic name).

  • $KNO_3$: Potassium nitrate.

• Anion Name Conversions:

  • Nitrogen $\rightarrow$ Nitride

  • Phosphorus $\rightarrow$ Phosphide

  • Oxygen $\rightarrow$ Oxide

  • Sulfur $\rightarrow$ Sulfide

  • Selenium $\rightarrow$ Selenide

  • Tellurium $\rightarrow$ Telluride

  • Fluorine $\rightarrow$ Fluoride

  • Chlorine $\rightarrow$ Chloride

  • Bromine $\rightarrow$ Bromide

  • Iodine $\rightarrow$ Iodide

• Stock System (Transition Metals): Because transition metals can have multiple charge states, Roman numerals must be used in the name to indicate the specific charge of the cation.

  • Example: Iron and Chlorine:

    1. $FeCl_2$: Iron is in a $+2$ state (indicated by the $2$ on chlorine). Name: Iron(II) chloride.

    2. $FeCl_3$: Iron is in a $+3$ state. Name: Iron(III) chloride.

  • Example: Chromium and Sulfur:

    1. $Cr_2S_3$: The $3$ on Sulfur came from Chromium ($Cr^{3+}$). Name: Chromium(III) sulfide.

Advanced Ionic Formula Examples

• Magnesium Nitride:

  • Anions: Magnesium ($Mg^{2+}$), Nitrogen ($N^{3-}$).

  • Interchange charges: $Mg_3N_2$.

  • Verification: $(3 \times (+2)) + (2 \times (-3)) = 0$.

• Copper(II) Nitrate:

  • Cation: Copper ($Cu^{2+}$). Anion: Nitrate ($NO_3^-$).

  • Formula: $Cu(NO_3)_2$.

• Potassium Dihydrogen Phosphate:

  • Cation: Potassium ($K^+$). Anion: Dihydrogen phosphate ($H_2PO_4^-$).

  • Formula: $KH_2PO_4$.

• Ammonium Chlorate:

  • Cation: Ammonium ($NH_4^+$). Anion: Chlorate ($ClO_3^-$).

  • Formula: $NH_4ClO_3$.

• Mercury(I) Nitrite:

  • Cation: Mercury(I) ($Hg_2^{2+}$). Anion: Nitrite ($NO_2^-$).

  • Formula: $Hg_2(NO_2)_2$.

• Cesium Sulfide:

  • Cation: Cesium ($Cs^+$). Anion: Sulfur ($S^{2-}$).

  • Formula: $Cs_2S$.

• Calcium Phosphate:

  • Cation: Calcium ($Ca^{2+}$). Anion: Phosphate ($PO_4^{3-}$).

  • Formula: $Ca_3(PO_4)_2$.

Molecular Compounds: Rules and Nomenclature

• Placement Rules:

  1. The element farthest to the left in a period on the periodic table is placed first in the formula.

  2. If elements are in the same group, the element closest to the bottom of the group is placed first.

• Prefixes for Molecular Compounds: Prefixes are used to indicate the number of atoms of each element present.

  • 1: Mono-

  • 2: Di-

  • 3: Tri-

  • 4: Tetra-

  • 5: Penta-

  • 6: Hexa-

  • 7: Hepta-

  • 8: Octa-

  • 9: Nona-

  • 10: Deca-

• Naming Rules:

  • The first element retains its full name.

  • The last element ends in -IDE.

  • If the first element has only one atom, "mono-" is typically omitted.

  • If a prefix is used for the first element, a prefix MUST be used for the second element, even if it is one (e.g., dinitrogen monoxide).

• Examples of Molecular Naming:

  • $HI$: Hydrogen iodide (single atoms, no prefixes needed).

  • $NF_3$: Nitrogen trifluoride.

  • $SO_2$: Sulfur dioxide.

  • $N_2Cl_4$: Dinitrogen tetrachloride.

  • $NO_2$: Nitrogen dioxide.

  • $N_2O$: Dinitrogen monoxide.

  • $SiCl_4$: Silicon tetrachloride.

  • $P_4O_{10}$: Tetraphosphorus decaoxide.

• Writing Formulas from Names:

  • Carbon disulfide: $CS_2$ (Di means $2$ Sulfurs).

  • Disilicon hexabromide: $Si_2Br_6$ (Di means $2$ Silicons, Hexa means $6$ Bromines).