periodic table: elements combining (chapter 3)

Bonding, Formulas, and Nomenclature — Notes

• Is essential distinction: ionic compounds form extended crystal lattices; formula units represent the smallest repeating ratio of ions in the lattice. For NaCl, the ratio is 1:1, so we write the formula as NaCl. In any real salt sample, you’ll have many such units, not a single Na and a single Cl.

• Empirical formula vs molecular formula:

  • Empirical formula = smallest whole-number ratio of atoms in a compound.

  • Molecular formula = actual number of each type of atom in a molecule.

  • For ionic compounds, we usually use empirical formulas or formula units; these often match the simplest ratio in the lattice.

  • For covalently bonded (molecular) compounds, the molecular formula shows the actual numbers in the molecule; sometimes the empirical and molecular formulas are the same (e.g.,
    $\text{H}_2\text{O}$), but not always.

• Electrostatic potential energy and Coulomb's law (conceptual):

  • The energy of interaction for two charges is proportional to the product of their charges divided by the distance between them:
    $E \propto \frac{q<em>1 q</em>2}{r}$

  • The stronger the charges (larger |q|), the stronger the interaction; smaller distances (r) allow charges to come closer, increasing the interaction.

  • In an ionic bond, a negatively charged ion (anion) and a positively charged ion (cation) attract more strongly as they approach, reach a minimum energy at an equilibrium separation, and then experience repulsion if forced closer due to nuclear repulsion.

  • The potential energy curve for ionic interactions shows a minimum at the bond distance; beyond this, repulsion rises as atoms get too close.

  • Note: This class discussion presents Coulomb’s law conceptually; you won't be required to memorize the equation at this stage.

• Ionic compounds vs covalent (molecular) compounds: quick distinctions

  • Ionic compounds: metal cations + nonmetal anions; form extended crystal lattices; empirical formula/ formula unit; examples include NaCl, MgO, CaS.

  • Covalent (molecular) compounds: nonmetals bonded by covalent (electron-sharing) bonds; form discrete molecules; molecular formulas describe the composition (e.g., $\mathrm{C}<em>2\mathrm{H}</em>6\mathrm{O}$ for ethanol).

  • Bonding curve analogy: for two nonmetal atoms at large separation, energy is effectively zero; as they approach, there is attraction due to shared electrons forming a covalent bond; a minimum energy occurs at the bond length, $r_{\text{bond}}$; pushing them closer leads to repulsion.

• Example molecules and formulas

  • Sodium chloride, NaCl: ionic compound; formula unit/empirical formula is NaCl; real samples contain many units.

  • Ethanol, $\mathrm{C}<em>2\mathrm{H}</em>6\mathrm{O}$: covalent molecular compound; molecular formula shows the actual ratio; empirical formula may be the same as molecular in this case.

  • Water, $\mathrm{H}_2\mathrm{O}$: an example where empirical and molecular formulas coincide.

  • Carbon allotropes: carbon exists as different forms (allotropes) such as diamond and graphite; both are elemental carbon but with different bonding structures.

  • Oxygen forms two common elemental diatomic molecules: $\mathrm{O}<em>2$ (oxygen gas) and $\mathrm{O}</em>3$ (ozone).

  • Diatomic molecules in general: $\mathrm{N}<em>2$, $\mathrm{O}</em>2$, $\mathrm{H}_2$, and the halogen diatomics working in nature.

  • Allotropy concept: a single element exists in more than one form with different properties (e.g., carbon as diamond vs graphite; oxygen as $\mathrm{O}<em>2$ vs $\mathrm{O}</em>3$).

  • Elemental molecules vs atomic elements: some elements exist as discrete atoms in nature (atomic elements, e.g., noble gases), while others exist as diatomic or polyatomic molecules (molecular elements, e.g., $\mathrm{O}<em>2$, $\mathrm{N}</em>2$).

  • Polyatomic ions: charged molecular species that can form ionic compounds (e.g., nitrate, $\mathrm{NO<em>3^-}$; ammonium, $\mathrm{NH</em>4^+}$).

• Classification of pure substances (overview)

  • Elements vs compounds

  • Elements can be:

  • Atomic elements (exist as single atoms, e.g., noble gases)

  • Molecular elements (exist as molecules, e.g., $\mathrm{O}<em>2$, $\mathrm{N}</em>2$, $\mathrm{H}_2$; diatomic halogens also fall here)

  • Compounds can be:

  • Molecular compounds (nonmetals bonded covalently; have molecular formulas)

  • Ionic compounds (metal + nonmetal; rely on ionic bonds; often described by empirical formulas)

  • Note on allotropes and elemental molecules: allotrope is a different form of the same element (e.g., carbon in diamond vs graphite). The terms $\mathrm{P}<em>4$ and $\mathrm{S}</em>8$ are examples of elemental molecules for phosphorus and sulfur.

• Naming binary ionic compounds (nomenclature rules, overview)

  • Rules (as discussed):

  • The first word is the cation (the metal, the element from the left side of the periodic table in the compound).

  • The second word is the anion name, with the suffix changed to -ide (indicating a negatively charged ion).

  • Examples:

    • Magnesium chloride: $\mathrm{MgCl_2}$ → Magnesium (cation) + chloride (an anion with -ide suffix)

    • Potassium bromide: $\mathrm{KBr}$ → Potassium + bromide

    • Calcium sulfide: $\mathrm{CaS}$ → Calcium + sulfide

  • Cross-method for balancing charges (simplified explanation):

  • To ensure overall neutrality, balance the charges of the ions by using the smallest whole-number subscripts that make the total charge zero.

  • Example: magnesium nitride

    • Ions: $\mathrm{Mg^{2+}}$ and $\mathrm{N^{3-}}$

    • Balance by cross-multiplication: the subscripts become $\text{Mg}<em>3\text{N}</em>2$ (because $2$ charges of Mg must balance $3$ charges of N; smallest integers that balance: 3 Mg2+ = +6 and 2 N3- = -6)

    • The balanced formula is $\mathrm{Mg<em>3N</em>2}$.

  • A general approach: use the magnitudes of the charges to determine the subscripts, then reduce if possible (common factors).

  • Common, but sometimes imperfect, rules encountered in class discussions (to be aware of):

  • Group 1 alkali metals are typically +1.

  • Group 2 alkaline earth metals are typically +2 (note: the transcript contains a misstatement here as part of its explanation; the correct charge is +2 for Group 2).

  • Nonmetals such as nitrogen typically take -3, oxygen typically -2, and halogens typically -1 (the transcript describes some inaccuracies here; the standard charges follow the periodic table and the octet rule).

  • The above charge assignments are the basis for predicting formulas of binary ionic compounds; always verify with proper oxidation state rules when applying to real problems.

• Practice problems (classification and basic naming in ionic context)

  • Task: Classify each substance as atomic element, molecular element, molecular compound, or ionic compound.

  • Example discussion points from the transcript:

  • Calcium oxide is an ionic compound (CaO) with Ca^2+ and O^2-; forms